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THIS book is intended for teachers who wish to emphasize the facts, laws, 
theories, and applications of chemistry. It is divided into two parts. Part I 
contains the text, together with exercises and problems. Part II contains the 
experiments. The text has been selected and arranged with special refer- 
ence to the needs of teachers as well as to the capacity of students. The 
experiments have been prepared to meet the needs of those schools in which 
the laboratory facilities are limited or the time for chemistry is short. 

The point of view differs from that in the author's " Experimental Chem- 
istry," but the spirit is the same. The two books are companion volumes, 
though of course they' can be used independently. The cordial reception 
given the " Experimental Chemistry " shows that many teachers are empha- 
sizing the experimental side of chemistry. These teachers will find Part I 
of the " Descriptive Chemistry " a serviceable companion book both in the 
laboratory and class room. It has been bound as a separate volume to meet 
such a use. 

Solutions of problems, answers to some of the exercises, and references to 
the literature have been put in a separate Teacher's Handbook. 

The manuscript has been read by Dr. William B. Schober, Lehigh Uni- 
versity, Bethlehem, Pennsylvania; Mr. Franklin T. Kurt, Chauncey Hall 
School, Boston, Massachusetts; and Mr. George M. Turner, Masten Park 
High School, Buffalo, New York. The chapters on theory were also read by 
Dr. Alexander Smith of the University of Chicago, and the chapters on 
carbon by Dr. James F. Norris of the Massachusetts Institute of Technology. 
The proof has been read by Dr. E. H. Kraus, High School, Syracuse, New 
York; Professor E. S. Babcock, Alfred University, Alfred, New York; and Mr. 
E. R. Whitney, High School, Binghamton, New York. The author is grateful 
to these teachers for their criticism, but he assumes all responsibility for any 
errors which may be detected. 

L. C. N. 






PART I. , f } 




III. HYDROGEN . . . .... . . . .23 






EQUATIONS .... -^^^^^- 100 







J Contents. 






XX. METALS .278 


XXII. COPPER SILVER GOLD . . . . . .301 



XXVI. TIN LEAD . . . .354 






APPENDIX ............ 437 





CHEMISTRY is a branch of natural science. It deals 
with the properties of matter, the changes which affect 
the composition of matter, with numerous laws and 
theories, and with the manufacture of a vast number of 
different substances indispensable to the welfare of man- 

Properties of Matter. Different substances are recog- 
nized and distinguished by their properties. Color, odor, 
taste, weight, and solubility are familiar properties ; but to 
these must be added behavior with heat, light, and electric- 
ity, and especially the action of different kinds of matter 
upon each other. 

Physical and Chemical Changes. Observation shows 
that the properties of matter can be changed. Sometimes 
the change is only temporary, as in the freezing of water, 
or in the melting of iron. Such changes are called physi- 
cal changes. But often the change is permanent, as in 
the burning of coal, or the digestion of food. Such 
changes are called chemical changes. In physical 
changes the original properties reappear after the cause 
of the change has been removed. But chemical changes 

2 Descriptive ^Chemistry. 

affect the essential nature of a substance. They are 
fundamental. Removal of the cause of a chemical change 
does not restore the original properties of the substance. 
Thus, coal is readily changed into ashes and invisible 
gases, but the ashes and gases do not reunite into coal 
after the heat has been removed. Another essential char- 
acteristic of chemical changes is the formation of one or 
more kinds of matter different from the original substance. 
Thus, water may be decomposed by electricity into two 
gases hydrogen and oxygen. This is a chemical change, 
because (i) the water has disappeared, its identity is lost, 
it has been permanently changed, and (2) other kinds of 
matter have been formed, which are totally unlike water. 
Chemistry is largely a study of chemical changes. 

The different changes which matter undergoes furnish 
a convenient basis for the classification of properties. 
Thus, we call physical properties those which accompany 
physical changes ; while chemical properties require a 
chemical change for their manifestation. Thus, the color, 
luster, specific gravity, melting point, and capacity to con- 
duct electricity are physical properties of copper; but it 
displays chemical properties when it is heated, or when 
acted upon by acids, sulphur, and other substances. 

Examples of simple physical changes are the formation of ice or steam 
from water, the electrification of a copper trolley wire, the production of 
colors in the sky, the magnetization of iron in a dynamo or magnet, and 
\the melting of iron in a foundry. Familiar chemical changes are the 
rusting of iron, the growth of plants, the burning of oil in a lamp, the 
decay of fruit, and the souring of milk. 

Chemical changes are often complex. In many in- 
stances they are caused by heat, and usually they produce 
heat. In general, the velocity of chemical change in- 
creases with rise of temperature. Light induces chemical 

Introduction. 3 

changes, as in growing plants and on photographic plates. 
Electricity is involved in many chemical changes, a vast 
industry having recently grown up in this field. Contact 
is necessary for chemical change, and many substances 
must be pressed together, intimately mixed, or dissolved 
before they will interact. 

Physical and chemical changes are closely related. 
They usually accompany each other, and are often insep- 
arable. If the essential change in a substance or sub- 
stances is chemical, then the substances are said to 
undergo chemical action. Very often the chemical action 
involves several substances. The substances are then said 
to interact or react, and the series of changes is called a 
reaction. Thus, when zinc is added to nitric acid, the 
chemical action which occurs is manifested by the forma- 
tion of a brown gas and the disappearance of the zinc. 
The zinc and acid interact, and tlie chemical changes can 
be classified as due to the reaction between zinc and nitric 

Classes of Chemical Action. There are four general 
kinds of chemical action, (i) Analysis or decomposition 
is the separation of matter into its components. Thus, 
heat decomposes wood, and the juices of our bodies de- 
compose food. (2) Synthesis or combination is the union 
of different kinds, or sometimes the same kind, of matter. 
For example, the gases, hydrogen and oxygen, may be 
made to unite and form water by passing an electric spark 
through them. (3) Substitution is the replacement of 
one kind of matter by another. When zinc is added to 
hydrochloric acid, the hydrogen leaves the acid, and zinc 
takes its place. (4) Sometimes parts of different sub- 
stances exchange places ; this kind of change is called 
metathesis or double decomposition. If silver nitrate is 

4 Descriptive Chemistry. 

added to hydrochloric acid, the silver and hydrogen ex- 
change places, forming silver chloride and nitric acid. 
These four kinds of chemical changes will be fully illus- 
trated and studied in the succeeding pages. 

Chemical Energy. We learn in physics that heat, 
light, and electricity are different forms of energy. They 
produce special changes. It is also possible to transform 
the different kinds of energy into each other. Thus, elec- 
tricity is generated from the heat liberated by burning 
coal, and electricity in turn may be transformed into light. 
In chemistry we study another kind of energy, called 
chemical energy, chemical attraction, or chemism. This 
is the immediate agent involved in chemical change. Com- 
bination and decomposition are due to its operation. 
Chemical energy may be transformed into light, electricity, 
and heat, and vice versa. Appreciable heat often accom- 
panies chemical changes, and we shall have many illustra- 
tions of the intimate relation between heat and chemical 
energy. Electricity is produced in an electric battery by 
chemical action. Light is one result of the chemical 
action called combustion or burning. In fact, every chemi- 
cal change is accompanied by an energy change of some 
kind, and in such transformations all the energy can be 
accounted for, none is lost or gained. 

Chemical energy is an essential factor in all chemical 
changes, but we know little or nothing of its nature. We 
can only study its results and its manner of action. 

Conservation of Matter. In chemical changes matter 
is not created or destroyed. It is often transformed, and 
apparently lost, but the total weight of the substances par- 
ticipating in any chemical change is always the same. 
The fact that matter is indestructible was first demon- 

Introduction. 5 

strated by the French chemist, Lavoisier (1743-1794), and 
countless observers have since shown that it is a funda- 
mental law of chemistry. The law is called the Law of 
the Conservation of Matter, and is often stated thus : - 

No weight is lost or gained in a chemical change. 

Chemical Elements. Study of the constitution of 
matter shows that some kinds can be decomposed into 
substances totally unlike the original matter. Water, for 
example, is easily decomposed into the gases, hydrogen 
and oxygen, which are entirely different from water. But 
it is impossible by any known process to obtain from 
some kinds of matter substances which have simpler prop- 
erties than the original substance. Thus, neither oxygen 
nor hydrogen can be decomposed by any known means. 
Iron and the familiar metals likewise cannot be divided 
chemically into two or more substances, nor can they be 
transformed into each other. They are fundamental sub- 
stances. We can add other substances to them, but we 
cannot get simpler substances from them, nor can we 
transform them into simpler substances. Iron contains 
nothing but iron. The substances which have such simple 
properties and at present defy decomposition and trans- 
formation are called the chemical elements. They are 
analogous to the letters of the alphabet, and by their vari- 
ous combinations make up the matter of the universe, some- 
what as letters form words. 

There are about eighty elements. Probably there are 
some undiscovered, but it is generally believed that the 
present number will not be largely increased. 

Each element is designated by a symbol, which is an 
abbreviation of its name. The following is an alphabeti- 

Descriptive Chemistry. 





Aluminium .... 


Lead .... 


Antimony .... 






Magnesium .... 













Nitrogen .... 















Chromium .... 
Cobalt . . 



Ae 1 








Strontium .... 













Of the above elements only eight are abundant in the 
earth's crust, as may be seen by a 



Iron . 
Potassium . 




7 .8l 



The atmosphere contains about 20 per cent of oxygen 
and 79 per cent of nitrogen in the free state. The ocean 
contains about 86 per cent of oxygen, 1 1 per cent of hydro- 
gen, and 2 per cent of chlorine in combined states. It is 
clear that the globe, as we know it, is made up of a very 
few elements. 

Many of the familiar metals are elements, e.g. lead, zinc, 
tin, copper, iron, gold, and silver. Other elements besides 
the metals are solids, such as sulphur, carbon, and phos- 
phorus ; two are liquid, viz. bromine and mercury ; while 
several are the common gases, oxygen, nitrogen, and hydro- 
gen. Many are important simply because they are com- 
bined with other elements, especially silicon, which is 
found in most rocks, and calcium, which is a component 
of limestone. 

The following is a 




















Ruthenium .... 



Gallium .... 


Scandium . 












Tellurium . . . . 
























Vanadium . 


Neon . 













Zirconium .... 



Descriptive Chemistry. 

Chemical Symbols are usually the first letter of the 
name of the element. Thus, O is the symbol of oxygen, 
H of hydrogen, N of nitrogen. Since several elements 
have the same initial letter, the symbol of some elements 
contains two letters. Thus, C represents carbon, while 
the symbol of calcium is Ca, of chlorine Cl, of chromium 
Cr, and of copper Cu. The symbols of several elements, 
especially the metals so long known, are derived from 
their Latin names, as may be seen from a 






































Symbols always begin with a capital, and are not followed by a 
period. They should be learned by actual use. Their significance 
will be explained in later chapters. 

Chemical Compounds. When elements unite with 
each other the product of the union is a chemical com- 
pound. The elements which make up a chemical com- 
pound are called components. Chemical compounds have 
three essential characteristics, (i) Their components are 
held together by chemical attraction. The hydrogen and 
oxygen, which are the components of water, cannot be 
separated unless their attraction for each other is over- 
come by heat, electricity, or some other agent. (2) In any 
given chemical compound the components are always in 

Introduction. g 

the same ratio. Thus, pure common salt, however pre- 
pared or wherever found, always contains 39.32 per cent 
of sodium and 60.68 per cent of chlorine. So also water 
always contains eight parts (by weight) of oxygen and one 
of hydrogen. Facts similar to these might be given cover- 
ing all cases examined. Such facts illustrate the general 
principle that chemical action proceeds according to laws. 
(3) In chemical compounds the identity of the components 
is lost. Thus, the red metal, copper, the yellow solid, 
sulphur, and the invisible gas, oxygen, are the components 
of the blue solid, copper sulphate. 

Chemical compounds must not be confused with mixtures. The 
parts of a mixture may vary in nature and in proportion ; they are also 
held together loosely, and may often be separated by some mechanical 
operation, as filtering or sifting. A mixture, too, often has properties 
similar to its parts. 


1. State three properties of (a) glass, (<) wood, (c} water, (W) paper, 
(e) air. 

2. Give three illustrations of (a) physical changes and (6) chemical 
changes occurring in everyday life. 

3. Are the following changes physical or chemical? (a} Burning 
of wood, () melting of butter, (c) freezing an ice-cream mixture, (d} 
weathering (i.e. decay) of granite, (e) tarnishing of brass and other 
metals, (/) formation of snow, (g) developing a photographic plate, (h) 
seasoning of wood, (/) formation of dew, (/) disappearance of a fog. 

4. What ai^ls and what retards chemical change? What often ac- 
companies it? 

5. What physical change accompanies (a} the burning of coal, (6) 
the action of an electric battery, (c) the burning of a match ? 

6. Give an illustration of the transformation of chemical energy into 
heat, light, or electricity. 

7. State the law of the conservation of matter. 

1 These exercises are intended for review work. 

io Descriptive Chemistry. 

8. (rt) Name five elements with which you are familiar. () Name 
the eight most abundant elements in the earth's crust in their order. 

9. What common metals are elements? 

10. How do elements and compounds essentially differ? Could you 
prepare (a) a compound from elements, (^) elements from a compound, 
and (c) elements from elements? 

11. Define (a) chemistry, () physical change, (c) chemical change, 
(d) chemical action, (^) analysis, (/") synthesis, (g) metathesis, (^) sub- 
stitution, (/) element, (/) compound, () mixture, (/) symbol. 

12. Review or learn the metric system (see Appendix, i). 


Perform the problems in the Appendix, i, 


OXYGEN has played an important part in the develop- 
ment of chemistry, and is an appropriate element with 
which to begin a systematic study of this science. 

Occurrence. Oxygen is the most abundant atid widely 
distributed of the elements. Mixed with nitrogen and a 
few other gases, it forms one fifth (by volume) of the 
atmosphere. Combined with hydrogen, it constitutes 
eight ninths (by weight) of water; combined with silicon 
and certain metals, it makes up nearly half of the earth's 
crust; while compounds of oxygen, carbon, and hydrogen 
form a large part of animal and vegetable matter. Starch, 
for example, which is a constituent of all plants, contains 
about 50 per cent oxygen. 

Preparation. Oxygen may be prepared from its com- 
pounds or from air. It was first prepared by decomposing 
a red compound of oxygen and mercury. When heated 
in a hard glass tube, this compound decomposes into 
oxygen and mercury ; the oxygen is collected over water 
in a pneumatic trough, and the mercury condenses as 
globules or a film on the upper part of the tube. This 
experiment is historically interesting, because it was first 
performed by Priestley, the discoverer of oxygen. 

The gas is often prepared by decomposing potassium 
chlorate a compound of oxygen, chlorine, and potassium. 
Heated to a rather high temperature, the potassium chlo- 

12 Descriptive Chemistry. 

rate passes through a series of changes ; as a final result, 
the oxygen is set free, and potassium chloride, a white 
solid, remains behind. 

Oxygen is most conveniently prepared by heating a 
mixture of potassium chlorate and manganese dioxide in a 
glass or metal vessel. The gas is liberated freely from 
this mixture at a lower temperature than when either 
compound is heated alone. 

The manganese dioxide may be recovered unchanged at the close of 
the experiment. It takes some part in the chemical changes, but just 
what is not definitely known. It has been suggested that the manganese 
dioxide combines at first with oxygen, thereby forming another coin- 
pound of manganese richer in oxygen than the dioxide, but so unstable 
that when heated it yields oxygen and manganese dioxide. 

Large quantities of oxygen may be prepared by heating a mixture of 
potassium chlorate and manganese dioxide in a copper or iron retort. 
Other commercial processes are used. In Erin's process, which is oper- 
ated largely in England, purified air is forced by a pump over barium 
oxide heated to 700 C., 1 thereby forming barium dioxide. The air sup- 
ply is then cut off, and the pressure in the retorts reduced by reversing 
the pump. This operation changes the barium dioxide into barium oxide 
and oxygen. The gas is drawn off into a reservoir. The process is 
then repeated. A kilogram of barium oxide yields about ten liters of 
oxygen at a single operation. 2 

Oxygen can be prepared from liquid air (see Liquid Air). By evapo- 
ration at the ordinary temperature and pressure, the nitrogen escapes 
from the liquid air more rapidly than the oxygen, leaving finally a liquid 
which is nearly pure oxygen. Unlimited quantities of oxygen may thus 
be cheaply prepared from the air. This method awaits development. 

Properties. Oxygen gas has no color, odor, or taste. 
It is slightly heavier than air. It is somewhat soluble in 

1 C. is the abbreviation of " centigrade," which is the name of the thermometer 
used in science. According to this thermometer water boils at 100 and freezes at 
o (see Appendix, 2). 

2 " Kilogram " and " liter" are denominations of the Metric System of Weights 
and Measures. This system should be learned or reviewed (see Appendix, i). 

Oxygen. 13 

water, but the presence of even a. small proportion in 
water is exceedingly important. Fish die in water con- 
taining no oxygen; and the oxygen absorbed by flowing 
water helps keep it free from organic matter. (See Decay, 

The density of oxygen gas is 1.105 (air = i). One hundred liters 
of water dissolve only about three liters of oxygen under ordinary 

The chemical activity of oxygen is its most striking 
property. It combines with all the other elements except 
fluorine, bromine, and the inert gases recently discovered 
in the atmosphere. With most of them the union is 
direct, and is often accompanied by light and heat, 
though the temperature at which combination occurs 
varies between wide limits. At the ordinary temperature 
it unites with phosphorus, as may be seen by the glow and 
fumes when the end of a match is rubbed, especially in a 
dark room. Metals, such as iron, lead, zinc, and copper, 
tarnish or rust easily, i.e. they combine with the oxygen 
of the air. The chemical activity of oxygen at high tem- 
peratures is readily shown by putting burning substances 
into it. All burn vividly in oxygen. 

When a glowing stick of wood is put into oxygen, the stick instantly 
bursts into a flame ; and if left in' the oxygen, the wood continues to 
burn brightly until the gas is exhausted. If glowing charcoal is put 
into oxygen, the charcoal burns violently, and throws off showers of 
sparks. Sulphur burns in air with a small, blue flame, but in oxygen 
the flame is much larger and brighter. The flame in both cases is 
accompanied by fumes which smell like a burning sulphur match. Iron 
wire does not burn in air, but if the end is coated with burning sulphur 
and then put into oxygen, the wire burns vividly, throwing off a shower 
of sparks ; when the flame has disappeared, a globule of red-hot iron is 
often seen on the end of the wire ; and sometimes the inside of the 
bottle is coated with a reddish powder, which is mainly a compound 

14 Descriptive Chemistry. 

of iron and oxygen. Iron and oxygen combine at a higher tempera- 
ture than do sulphur and oxygen, so sulphur is used to set fire to the 
iron. On the other hand, if lighted magnesium is put into oxygen, the 
burning metal instantly becomes surrounded with a dazzling flame, and 
burns rapidly to a white powder, thus showing that the temperature at 
which it combines with oxygen is much lower than that required by iron. 

Oxidation. When sulphur, iron, magnesium, and car- 
bon (in wood and charcoal), and other elements burn in 
oxygen, they combine with it. This chemical change is 
called oxidation. 

The fact that oxidation is merely a combining with oxygen may be 
easily verified. It has been repeatedly shown that oxygen is one con- 
stituent of all the products formed by burning substances in that gas. 
Thus, carbon forms an invisible gas called carbon dioxide, which is a 
compound of carbon and oxygen. Similarly, sulphur, iron, and magne- 
sium form compounds of these elements and oxygen. These facts may 
be further verified by a simple experiment. If mercury is heated, it 
gains in weight, and red particles collect on its surface ; but if it is pro- 
tected from the air by some coating and then heated, there is no gain 
in weight and no evidence of the red product. Therefore, when the 
exposed mercury is heated, something from the air must be added to it. 
Now, if the red substance is collected and heated in a glass tube, mercury 
and oxygen are the only products. Hence, the exposed mercury, when 
heated, must have combined with the oxygen of the air. 

Oxidation is not always rapid enough to produce light 
and appreciable heat. Iron and other metals rust, and 
wood decays slowly, but both processes are mainly oxida- 
tion. Sometimes oxidation develops considerable heat. 
Thus, oily rags, piles of hay, and heaps of coal often take 
fire unexpectedly because of the continued oxidation. Such 
oxidation is often called spontaneous combustion. 

Substances which give up oxygen readily are called 
oxidizing agents. Potassium chlorate is used in fireworks 
for this purpose, and potassium nitrate acts similarly in 
gunpowder. In the process of oxidation, oxidizing agents 

Oxygen. 15 

lose oxygen, and are said to undergo reduction a process 
which will be more fully described in the next chapter. 

Oxides are formed when oxygen combines with other 
elements. There are many oxides, and their names express 
in a general way their composition. Oxides of different 
elements are distinguished by placing the name of the ele- 
ment (or a slight modification of it) before the word oxide, 
e.g. magnesium oxide, lead oxide, zinc oxide. Sometimes 
di-, or a similar numerical syllable, is prefixed to the word 
oxide, e.g. carbon dioxide, manganese dioxide, sulphur 
trioxide, phosphorus pentoxide. The significance of the 
prefix is explained in Chapter VII. 

Combustion, in a narrow sense, is rapid oxidation, which 
is always accompanied by light and heat. Popularly, com- 
bustion means fire or burning, and substances which burn 
easily are called combustible. Oxygen is essential to ordi- 
nary combustion, and is often called a supporter of com- 
bustion. Exclude air from a fire, and the fire goes out. 
When coal or wood burns, the carbon (of which they 
largely consist) unites with the oxygen of the air, forming 
thereby the invisible gas carbon dioxide,, and the chemical 
change is manifested by heat and light. /Chemically speak- 
ing, a substance burning in the air is Uniting rapidly with 
oxygen. But since the air is about one fifth oxygen and 
four fifths nitrogen, a gas which does not support com- 
bustion, it follows that combustion is more vigorous in 
oxygen than in air. 

The correct explanation of fire, burning, and combustion was first 
made by Lavoisier (1743-1794). For many years chemists had be- 
lieved that all combustible substances contained a principle called 
phlogiston, and that when a substance burned, phlogiston escaped. 
Very combustible substances were thought to contain much phlogiston, 
and incombustible substances no phlogiston. This theory of combus- 

1 6 Descriptive Chemistry. 

tion was proposed by Becher (1635-1682) and advanced by Stahl 
(1660-1734). Many famous chemists Priestley, Scheele, and Caven- 
dish supported it. Lavoisier, in 1775, proved by his own and others 1 
experiments, that phlogiston did not exist, and that combustion is a 
process of combination with " a certain substance contained in the air." 
Soon after he identified this substance as oxygen. The theory of 
phlogiston, in spite of its falsity, exerted a wholesome influence on the 
development of chemistry. 

Combustion, in a broad sense, is not necessarily oxida- 
tion, but chemical action which develops enough energy 
to produce light and heat. This broader meaning will be 
discussed later. 

Relation of Oxygen to Life. Oxygen is essential to all 
forms of animal and plant life. If an animal or a plant is 
deprived of air, it dies. By respiration air is drawn into 
the lungs and there it gives up part of its oxygen to the 
blood. This oxygen, which is distributed to all parts of 
the body by trie blood, oxidizes food and the tissues of the 
body. As a result of this oxidation new tissue is built up 
and waste products are formed. One of these waste prod- 
ucts is carbon dioxide gas, which with other gases is 
exhaled from the lungs. The blood during its circulation 
turns dark red, owing to the loss of oxygen ; and when this 
dark red blood reaches the lungs, it receives a fresh supply 
of oxygen which turns it bright red, thus preparing it for 
another journey through the body. Food must be oxidized 
before it can be taken up by the body, and by this oxida- 
tion the carbonaceous matter of the body is slowly burned 
to carbon dioxide. It is this slow oxidation which keeps 
the body warm. The human body resembles a steam 
engine. In^each, the oxYggiLjQ-the-air.lielps_burn fuel ^ 
largely composed of carbon. In the engine, the products 
es"cape through a chimney and the heat produced is used 

Oxygen. 17 

to form steam which moves parts of the machine ; in the 
body, the products escape mainly through the lungs and 
the heat keeps the body at a temperature at which it can 
best perform its functions. 

It was formerly believed that breathing pure oxygen would produce 
too rapid oxidation in the body and burn up the tissue faster than it 
could be made. But recent study shows that with proper precautions 
oxygen may be breathed by a healthy person without producing any 
harmful effect. The blood apparently absorbs a maximum quantity of 
oxygen, whether supplied from air or from the pure gas. Oxygen is 
often administered to a person who has been suffocated, or to one who 
is unable to inhale enough air, as in cases of croup, asthma, or extreme 
weakness. It is sometimes used to sustain life where air is impure 
or rare, as in diving bells and submarine boats, and during balloon 
ascensions to a great height. 

Decay is in part oxidation. The oxygen of the air 
together with water vapor acts upon animal and vegetable 
matter and slowly burns it up. The decomposition is often 
begun and hastened by bacteria. The products of decay 
are numerous, carbon dioxide being one. The oxygen 
dissolved by water assists in the decay of the impurities 
constantly flowing into rivers. Similarly, it oxidizes in- 
jurious vapors and matter in the air, literally burning them 
up, just as it burns wood in a stove. Hence, running 
water is more likely to be cleaner than standing or stagnant 
water, and the air in the open country or at the seashore 
purer than in the crowded city. 

Uses of Oxygen. Oxygen for commercial use is stored under 
pressure in strong iron cylinders. The pure gas has limited use, since 
air, although it contains about 80 per cent of the inert gas nitrogen, 
may usually be used in place of oxygen, A mixture of oxygen and 
hydrogen burned in a suitable apparatus produces an intensely hot 
flame, which is sometimes used to melt refractory metals and to produce 
the calcium light (see Oxyhydrogen Blowpipe). 

1 8 Descriptive Chemistry. 

Liquid Oxygen. All gases at a low temperature and 
under great pressure may be condensed to liquids, and 
even to solids. Under these conditions oxygen becomes 
first a pale blue liquid and finally a whitish solid. A small 
quantity was first obtained in 1877, but now it is prepared 
by the gallon. It is magnetic, and when a strong electro- 
magnet is held near its surface, the liquid suddenly "leaps 
up to the poles and remains there permanently attached 
until it evaporates." 

Under the normal pressure (760 mm.) 1 liquid oxygen boils at 
181.4 C., and at this temperature its specific gravity is 1. 124 (water i). 

Discovery of Oxygen. Oxygen was discovered on 
August i, 1774, by Priestley (1733-1804). He prepared it 
by focusing the sun's rays upon the red mercury oxide by 
means of " a burning lens of twelve inches' focal distance." 
It was independently discovered by Scheele (1742-1786), a 
Swedish chemist, about the same time. 

Priestley called the gas dephlogisticated air, because he regarded it 
as " devoid of phlogiston." Scheele called it empyreal air, i.e. fire 
air or fire-supporting air, because it assisted combustion. Lavoisier, in 
1778, gave it the name oxygen (from the Greek oxus, acid, and^w, the 
root of a verb meaning to produce), because he believed from his 
experiments that oxygen was necessary for the production of acids a 
view now known to be incorrect. 

Weight of a Liter of Oxygen. The volume occupied 
by a gas depends upon the pressure and temperature to 
which it is subjected. The volume expands with rise of 
temperature or with lowering of pressure, but contracts 
with fall of temperature or with increase of pressure. In 
general, if we cool a gas or subject it to a pressure, it 
shrinks, and if we heat a gas or decrease the pressure 

1 This expression means the normal or standard pressure of the atmosphere as 
recorded by the barometer (see Chapter VI). 

Oxygen. 19 

it is under, it expands. Gas volumes, to be correctly 
compared, must therefore be at the same temperature and 
pressure. The normal or standard temperature is zero 
degrees on the centigrade thermometer, or briefly o C. 
The normal or standard pressure is the pressure of the 
atmosphere indicated by the barometer when the mercury is 
760 millimeters high, or briefly 760 mm. Under these 
conditions, which are called standard conditions, a liter of 
dry oxygen weighs 1.43 gm. 

It is not usually convenient to measure gases at o C. 
and 760 mm. So if their volumes are to be studied and 
compared, it is customary to reduce the observed volume 
to the volume it would occupy under standard conditions. 
This reduction is accomplished by applying two laws the 
Law of Charles and the Law of Boyle. 

Law of Charles. It has been found by experiment that under con- 
stant pressure all gases expand or contract equally for equal changes 
of temperature. More explicitly, a gas expands or contracts ^ of its 
volume at o C. for every degree through which it is heated or cooled. 
This means that 273 volumes at p become 274 at i, 275 at 2, 280 at 
7, 272 at i, 270 at 3, or 273 + 1 volumes at / (i e. at any tem- 
perature). This law is not absolutely correct, but its variations from 
the truth are slight. 

Suppose we have 10 1. of oxygen at o C., and we wish to know 
the volume it would occupy at 15 C. The problem is easily solved by 
stating it as a proportion, thus 

273:273+ I5::lo:.r. - 

The value of .ris the volume required. Conversely, in reducing 10 vol- 
umes at 15 C. to the volume occupied at o C., the proportion is 

273 + I5:27 3 ::io:;r. 

If the given temperature is below o, the number of degrees is subtracted 
from 273. 

Law of Boyle. It has also been found by experiment that under 
constant temperature the volume of a gas is inversely proportional to 

lo Descriptive Chemistry. 

the pressure. This is Boyle's law. It means that doubling the pres- 
sure halves the volume, and vice "versa. Like the above law, this law is 
only approximately correct. 

Suppose we have 10 1. of oxygen at 760 mm., and we wish to know 
the volume it would occupy at 775 mm. According to the law, the 
proportion expressing the relation is 

760:775::^: 10. 

The value of x\<& the required volume. Conversely, if we have 10 1. at 
775 mm., and wish to know its standard volume, the proportion is 

It is convenient to notice that the proportion is stated so that the 
extremes (or means) are the original pressure and volume. In other 
words, one pressure multiplied by its volume equals the other pressure 
multiplied by its volume, or 

P\P\\V\ V. 

Hence, the proportion is applicable to values not necessarily includ- 
ing 760. 


i . What is the symbol of oxygen ? 

2. How is oxygen prepared (a) in the laboratory, and () commer- 
cially ? 

3. Name several compounds from which oxygen can be prepared. 

4. Summarize the properties of oxygen. What is its most charac- 
teristic property ? 

5. If air contains something besides oxygen, what must be the gen- 
eral properties of this other ingredient ? 

6. Define and illustrate (a) oxidation, (ft) oxide, (c) combustion, 
(//) oxidizing agent. 

7. What elements were mentioned in studying oxygen ? What 
compounds ? 

8. What general chemical change is involved in burning ? What 
class of chemical changes is illustrated by (a} preparation of oxygen 
from mercuric oxide, (b} burning of sulphur in oxygen ? 

9. Give a brief account of Priestley, Scheele, and Lavoisier (see 
Appendix, 4). 

Oxygen. 21 

10. What chemical part does oxygen 'take in (a) respiration, (#) de- 
cay, (c) combustion, () oxidation ? 

11. State and illustrate (a) Charles's law and () Boyle's law. 

12. Give a brief account of Boyle and of Charles. 


1. Potassium chlorate contains about 39 per cent of oxygen. How 
many grams of oxygen can be prepared from (rt) 100 gm., (^) 250 gm., 
and (c) 725 gm. of potassium chlorate ? 

2. What approximate weight of oxygen can be prepared from 100 
gm. of potassium chlorate containing 12 per cent of impurity ? 

3. What is the weight of (a) 10 1. of oxygen, (b) 75-!., (c) 500 
cc., (d) 750 cc., 0) 4!.? 

4. A room 25 m. long, 17 wide, and 15 high is filled with oxygen. 
What weight of gas does it contain ? (A liter of oxygen weighs 

1-43 gm.) 

5. Reduce the following volumes to the volume occupied at o C. : 
(a) 173 cc. at 12 C., (b) 466 cc. at 14 C., (c) 706 cc. at 15 C., (d) 
25 cc. at 27 C. 

6. A volume of gas at o C. measures 1500 cc. What is its volume 
at (a} 15 C., (d) 50 C., (0 100 C., (d} 300 C. ? 

7. If 500 cc. of gas at 27 C. are cooled to 5C., what is the new 
volume ? 

8. Reduce the following volumes to the volume occupied at 
760 mm. : (a) 200 cc. at 740 mm., (b) 25 cc. at 780 mm., (c) 467 cc. 
at 756 mm. Ans. (a) 1947? (^) 25.65, (c) 464-54- 

9. A gas measures 1000 cc. at 770 mm. What is its volume at 
530 mm.? 

10. Reduce the following to standard conditions: (a) 147 cc. at 
570 mm. and 136.5 C., (b} 320 cc. at 950 mm. and 9iC, (c) 480 cc. 
at 380 mm. and 68.25C, (d) 25 cc. at 780 mm. and 27 C., (*) 14 cc. 
at 763 mm. and iiC. 

Ozone is a gas related to oxygen, though its properties differ. It is 
formed when electric sparks pass through the air, and is therefore pro- 
duced when electrical machines are in operation and during thunder 
storms. Slow oxidation, especially of moist phosphorus, produces 
ozone. Indeed, its formation accompanies several chemical changes. 

22 Descriptive Chemistry. 

such as the burning of hydrogen and of certain resins, and the decom- 
position of water by electricity. 

Ozone has a peculiar odor, suggesting burning sulphur. The name 
ozone signifies smell. It is active chemically, tarnishing metals, bleach- 
ing colored vegetable substances, deodorizing foul animal matter, and 
corroding such substances as cork and rubber. It is sometimes used as a 
disinfectant, though other oxidizing agents are more convenient. When 
heated to 250 C., or higher, it is wholly changed into oxygen. Ozone, 
therefore, contains nothing but oxygen. When oxygen is changed into 
ozone, it is found that three volumes of oxygen yield two volumes of 
ozone ; and, conversely, the two volumes of ozone, when heated, become 
three volumes of oxygen. Hence, volume for volume, ozone is 1.5 times 
heavier than oxygen. For this reason ozone is sometimes called "con- 
centrated oxygen," or "an oxide of oxygen." Its theoretical relation to 
oxygen will be subsequently discussed. 

The atmosphere usually contains a small proportion of ozone, prob- 
ably not more than one volume in 700,000 volumes of air. It is more 
abundant in the open country and at the seashore than in cities. 


Occurrence. Free hydrogen is present in the gases 

petroleum wells, and natural 

gas openings. Artificial illuminating gas contains consid- ^ 
erable hydrogen. It is^T product of fermentation and *"* 
decay, and according to recent observations a very small 
quantity is present in the atmosphere of the earth. Enor- * 
mous quantities of free hydrogen exist in tne atmqsrjhere 7 
of the sun, and during an eclipse of the sun gigantic 
streams of burning hydrogen may be seen shooting out 
from the sun's disk thousands of miles into space. Other 
heavenly bodies which are self-luminous, like the star Sirius 
and the nebulae, contain free hydrogen. The spectroscope 
has revealed its presence in these distant bodies. Meteor-*? 
ites^ which come from regions far beyond our earth, often 
contain free hydrogen. 

Cojnbinedji^drogen is abundant and widely distributed. * 
It forms one ninthly weight^ of water. Most animal and f ^ 
vegetable matter contains hydrogen. It is also an essential 1 1 
component nf_a1J_gHHs Combined with carbon, it forms 
many gases and liquids called hydrocarbons, which are con- ' * 
stituents of illuminating gas, kerosene, and naphtha. Com- 
bined with carbon and Oxygen, It forms many vegetable > C> 
compounds, such as sugar, starch, parser, wood, and numer- 
ous artificia'1 products. With nitrogen it forms the familiar ^ 
compound, ammonia ; and with sulphur, the bad-smelling gas, ^ 
hydrogen sufphTdeTwhich occurs in many sulphur springs. ' 


24 Descriptive Chemistry. 

Preparation. Hydrogen, like oxygen, is prepared from 
its compounds. In the laboratory this is easily accom- 
plished by allowing a metal and an acid to interact. The 
metals usually employed are zinc, iron, or magnesium, and 
the acids are dilute sulphuric acid or hydrochloric acid. 
The hydrogen comes from the acid and bubbles through 
the liquid, when the acid and metal are put into a test tube 
or flask. On a large scale hydrogen is prepared in a genera- 
tor, which consists of a glass vessel provided with a delivery 
tube arranged to collect the gas over water in a pneumatic 
trough. No flame should be near during the performance 
of this experiment, because mixtures of air and hydrogen 
explode violently when ignited. The interaction of zinc 
and sulphuric acid produces, besides hydrogen, a compound 
called zinc sulphate. This remains in the generator in 
solution, and if the solution is allowed to evaperate, the 
zinc sulphate separates as transparent crystals, which soon 
turn white in the air. Hydrogen may be obtained from 
water by allowing tiie_Jiilal-ao^lijLir^^ to interact. 

If a small piece of sodium is dropped upon cold water, the sodium 
melts into a shining globule, which spins about rapidly on the water 
with a hissing sound, and finally disappears with a slight explosion. 
But when the sodium is wrapped in a piece of tea lead pierced with a 
few holes and then dropped beneath the shelf of a pneumatic trough 
filled with water, the action proceeds smoothly. Hydrogen gas rises 
and displaces the water from a test tube or bottle supported over the 
hole in the shelf. The nature of the chemical change which attends 
the liberation of hydrogen from water will be explained later (Chap- 
ter V). 

Hydrogen, together with oxygen, is liberated from water 
by passing a current of electricity througlTwafer containing 
a little sulphuric acid (see Chapter V). 

Hydrogen may also be prepared by passing steam the 
gaseous form of water over heated metals. 

Hydrogen. 25 

This experiment was first performed by Lavoisier, in 1783, while he 
was studying the composition of water. He -passed steam through a 
red-hot gun barrel containing bits of iron. The oxygen of the steam 
combined with the iron, and the hydrogen escaped from the tube. Since 
Lavoisier was studying the composition of water, and not the properties 
of hydrogen, he naturally thought of this gas as essential for forming 
water. So he says in his notes, " No name appears to us more suitable 
than that of hydrogen, that is to say, 'generative principle of water.'" 
Apart from historical interest, this experiment has commercial value. 
If steam is passed over red-hot coal (instead of iron), producer gas 
is formed. This is a mixture consisting largely of hydrogen, which is 
used as a source of heat in making steel and glass. If oil vapor is 
added to this mixture, water gas is formed. This is an illuminating 
gas like ordinary illuminating gas, and is used in many cities (see 
Water Gas). 

Physical Properties. Hydrogen has no taste or color. 
The pure gas has no odor, though hydrogen as ordinarily 
prepared has a disagreeable odor, due mainly to impurities 
in the metals used. Most of these impurities may be re- 
moved by passing the gas through a solution of potassium 
permanganate. Hydrogen is-the lightest known substance. 
One liter of dry hydrogen at o C. and 760 mm. weighs 
only 0.0896 gm. Volume for volume, air is about 14.4 
times, oxygen 16 times, and water 11,000 times heavier 
than hydrogen. 

The extreme lightness of hydrogen may be easily shown, (i) If a 
wide-mouth bottle of the gas 
is left uncovered two or three 
minutes and a lighted match 
then dropped in, the match 
will continue to burn. If 
hydrogen had been present, 
the flame would have caused 
it to combine with the oxy- 

gen of the air with a loud FlG le l pouring hydroge n. 

explosion. (2) If a bottle of 
hydrogen is held beneath a bottle of air as shown in Figure i, the gases 

26 Descriptive Chemistry. 

soon exchange places, the hydrogen, owing to its lightness, rising into 
the upper bottle. Its presence there may be readily shown by dropping 
a lighted match into this bottle ; if the experiment has been well done, 
the hydrogen will burn, but in most cases the loud explosion shows 
that only a part of the hydrogen has been poured upward. A lighted 
match dropped into the other bottle reveals only air. (3) If a small 
collodion, or rubber, balloon is filled with hydrogen and then released, 
it will rise rapidly into the air. Hydrogen, because of its lightness, is 
sometimes used to fill large balloons, but ordinary illuminating gas is 
usually employed. 

Hydrogen is the standard for reckoning the density of gases. Thus, 
since a liter of oxygen weighs 1.43 gm., its density is found by the 
proportion:- Q ^ . , ^ . . , . ^ . ^ l6 

Hydrogen is not very soluble in water, but it is absorbed 
by several metals, especially the rare metal palladium. 
This property of absorbing gases is called occlusion. 

Only about 1.84 1. of hydrogen at 760 mm. pressure dissolve in 
100 1. of water at 20 C. Palladium absorbs from 370 to 960 times 
its own volume of hydrogen, according to the conditions of the experi- 
ment. Platinum and iron act similarly, though to a less degree. Illu- 
minating gas, which contains considerable hydrogen, is also absorbed 
by metals. And since heat is developed by occlusion, the illuminating 
gas may be lighted by the heated metal upon which it flows. A self- 
lighting gas burner acts on this principle. The act of occlusion is 
partly chemical and partly physical. 

Hydrogen illustrates diffusion; i.e. it readily passes 
through porous substances and completely mixes with 
other gases without stirring or agitating. 

It penetrates unglazed earthenware, paper, and heated metals, espe- 
cially platinum. Hydrogen has the highest rate of diffusion, because 
its density is the lowest. The rate of diffusion of a gas is inversely 
proportional to the square root of the density. Thus, the rate of diffu- 
sion of hydrogen is four times that of oxygen, since the density of oxy- 
gen is sixteen times that of hydrogen. We are largely indebted for 
our knowledge of diffusion to the English chemist, Thomas Graham 

Hydrogen. 27 

Hydrogen is not poisonous if pure. It does not sup- 
port life, but a little may be breathed without danger. 
When the lungs are filled with it the voice becomes very 
shrill and thin. 

Chemical Conduct. Hydrogen burns in the air and 
in oxygen with an almost invisible but very hot flame. 
Water is the product of its 
combustion. These facts may 
be verified by the apparatus 
shown in Figure 2. The hydro- 
gen, which is generated from 
zinc and hydrochloric acid in 
the flask, passes through the 
U-tube filled with calcium 
chloride (to remove the mois- 
ture), and is lighted at the tip 
after it has driven all the air from the apparatus. 1 A 
platinum or copper wire held in the flame instantly becomes 
red-hot. If a small, dry, cold bottle is held over the flame, 
moisture is deposited inside the bottle. 

The film of water often noticed on the bottom of a vessel placed 
over a lighted gas range or a Bunsen burner is formed by the burning 
hydrogen and hydrogen compounds of the illuminating gas. Similarly, 
water often drops from the top of the oven of a lighted gas range. Or- 
ganic substances containing hydrogen, such as wood and paper, when 
burned, yield water as one of their products. 

The fact that the only product of burning hydrogen is water was first 
shown in 1783 by Cavendish (1730-1810). Lavoisier in the same year 
verified this fact and utilized it to explain the composition of water. 

The temperature of the hydrogen flame is very high. 
More heat is produced by burning hydrogen in oxygen 

FIG. 2. Apparatus for burning 
hydrogen. . 

1 This experiment is dangerous. The precautions to be observed can be 
found on pages 48-49 in the author's " Experimental Chemistry." 

28 Descriptive Chemistry. 

than by burning the same weight of any other substance 
(see Chapter X). 

Hydrogen burns in chlorine gas. The flame is bluish white, not 
very hot, and the product is hydrochloric acid gas a compound of 
hydrogen and chlorine. This burning of hydrogen in chlorine illus- 
trates the broader use of the word combustion, since no oxygen is 

Hydrogen does not support combustion, as the term is 
usually used. This fact is illustrated by putting a lighted 
taper into an inverted bottle of hydrogen. The taper 
ignites the hydrogen, which burns at the mouth of the 
bottle. The taper does not burn inside the bottle, but when 
it is slowly withdrawn through the burning hydrogen it is 
relighted. Hence, hydrogen burns, but does not support 

A mixture of hydrogen and air explodes violently when 
ignited. Therefore, the air should be fully expelled from 
the apparatus in which hydrogen is being generated before 
the gas is collected, and no flames, large or small, should be 
near. Neglect of these precautions has caused serious 

Hydrogen not only combines energetically with frea 
oxygen, but it withdraws oxygen from compounds. As 
stated before, this chemical removal of oxygen is called 
reduction. Hydrogen is a vigorous reducing agent. 

The Oxyhydrogen Blowpipe utilizes the intense heat pro- 
duced by burning a mixture of hydrogen and oxygen. The 

apparatus (Fig. 3) con- 
sists of two pointed metal 
tubes. The inner and 
smaller one is for the 
Blowpipe tip. oxyg e n , an d the outer 

and larger one for the hydrogen. Their pointed ends are 

Hydrogen. 29 

close together, and the two gases mix as they are forced 
out of these small openings by the pressure maintained in 
the storage tanks. Sometimes the tubes are separated, 
but the gases flow from a similar opening. The hydrogen 
is first turned on and lighted at the pointed opening ; then 
the oxygen is turned on and the flow gradually regulated 
until the flame is the desired size, usually thin, straight, 
and as long as the apparatus requires. There is no danger 
in using the blowpipe, provided it does not leak and the 
pressure is properly regulated by the stopcocks. In the 
hot flame, some metals, like silver, turn to vapor ; some, 
like iron, burn brilliantly ; while others, like platinum, melt. 
When the flame strikes against a piece of lime of other sub- 
stance difficult to melt, the lime becomes intensely bright. 
Thus used, it is called the lime, calcium, or Drummond 
light and is often employed in operating the stereopticon. 

The blast lamp is a modification of the oxyhydrogen blowpipe. The 
apparatus (Fig. 4) consists of two tubes, an inner one for air and an 
outer one for illuminating gas. The air, 
which is forced through the apparatus by 
a bellows, provides oxygen, and the illumi- 
nating gas contains hydrogen and other 
combustible gases. The mixture burns at 
the opening of the tubes with a colorless 
or bluish flame, which is hotter than the 
Bunsen flame the usual source of heat for 
chemical experiments. The shape of the 
flame is easily regulated by stopcocks. 

Liquid Hydrogen is a colorless, trans- 
parent liquid produced bv subjecting the 

FIG. 4. Blast lamp, 
gas to great pressure and low temperature. 

It was first produced in 1898 by Dewar. The temperature used was 
205 C., and the pressure was 180 atmospheres (i.e. 180 times 760 
mm.). At the ordinary pressure it boils at 238 C. Under reduced 
pressure and at 256 C. it becomes "a white mass of solidified foam." 

jo Descriptive Chemistry. 

Discovery of Hydrogen. Paracelsus in the sixteenth century ob- 
tained hydrogen by the interaction of acids and metals. It was iden- 
tified as an element in 1766 by Cavendish, who called it inflammable 
air. The name hydrogen, given to it by Lavoisier, in 1783, is derived 
from the Greek words hudor, water, and gen, the root of a verb mean- 
ing to produce. 


1. What is the symbol of hydrogen ? 

2. What familiar compounds contain hydrogen? 

3. How is hydrogen prepared in the laboratory? Describe other 
methods of preparation. 

4. Summarize the properties of hydrogen. What is its most char- 
acteristic property ? 

5. Why is there danger of an explosion in generating hydrogen? 
How may the danger be avoided ? 

6. What is the weight of a liter of dry hydrogen? How many 
times heavier than a liter of hydrogen is one of air ? 

7. Define and illustrate (a) occlusion and (b} diffusion of gases. 

8. What chemical change occurs when hydrogen burns in air ? 

9. Is water an oxide ? Why ? 

10. How does the heat of the hydrogen flame compare with its 
luminosity ? 

n. Define (#) reduction and () reducing agent. Name a reduc- 
ing agent. 

12. Describe () the compound blowpipe and (&) the blast lamp, 
and state the use of each. 

13. Summarize briefly the discovery of hydrogen. Give a short 
account of Cavendish. Why and by whom was hydrogen so named ? 

14. What class of chemical changes is illustrated by () the prepara- 
tion of hydrogen from zinc and sulphuric acid, (<) the burning of 
hydrogen in air ? 


1. How many times heavier than a liter of hydrogen is a liter of 
oxygen, both being dry and under standard conditions ? 

2. What is the weight of (a) 500 cc. of dry hydrogen gas at o C. 
and 760 mm. ? (b) Of 1800 cc. ? (V) Of 9 1. ? 

3. The standard pressure at which a gas is measured is 760 mm. 
Express the same in inches. 


WATER is worthy of extensive study because of its 
importance in the animal, vegetable, and mineral king- 
doms, its peculiar properties, and its numberless uses. 

Occurrence in Nature. Water, in the form of vapor, 
is always present in the atmosphere. Evaporation is con- 
stantly taking place from the surface of the ocean, from 
the moist earth, from the bodies of animals, and from 
plants. This vapor is continually condensing, and appears 
as clouds, mist, fog, rain, snow, hail, dew, and frost. 

The proportion of water vapor in the atmosphere varies between wide 
limits, the amount present being largely influenced by the temperature. 
It has been found, however, that 1000 volumes of ordinary air contain 
about 14 volumes of water vapor. The total amount of vapor in the atmos- 
phere is beyond comprehension. 

In the liquid state water occurs in vast quantities. 
About three fourths of the surface of the globe is covered 
with water. Soil and porous rocks hold considerable 
quantities, and plants and animals contain a large pro- 
portion. Many substances which are apparently dry really 
contain a large proportion of water. Thus, in a ton of 
clover hay there are upwards of 200 Ib. of water, and a 
ton of salt hay, which is usually very dry, contains about 
100 Ib. 

Many common foods are largely water, as may be seen 
by the following 


Descriptive Chemistry. 






Cod .... 



Beef . . 






7Q 2 

Strawberries . . 

QO A. 

Ecrorg . 

/y * 

Watermelon .... 

y w "4- 







Cheese ....... 

28 to 72 



White bread .... 


The human body is nearly 70 per cent water, and during a 
year the average man drinks about half a ton. 

Water in the form of ice permanently covers the coldest 
parts of the surface of the earth, e.g. the polar regions and 
the summits of high mountains. A rough estimate of the 
total weight of ice on the earth's surface is 6,373,000,0x30 
millions of metric tons. 1 

Functions of Water in Nature. Since water is the 
only liquid occurring in large quantities on the earth's sur- 
face, it is the great agent of erosion. It cuts away the 
earth's crust, and transports the material from higher to 
lower levels, or washes it into the ocean. Together with 
carbon dioxide gas it decomposes the rocks, changing them 
into clay, sand, and substances which make the soil pro- 
ductive. Its cycle of changes from liquid to vapor and 
vapor to liquid exerts a marked influence on the distribu- 
tion of heat and moisture upon the earth's surface, i.e. on 

It dissolves many solids and gases and is constantly re- 
moving from the rocks and soil their soluble constituents, 

1 A metric ton contains 2204.6 pounds. 

Properties of Water. 33 

some of which serve for the nutrition of plants, though the 
larger part passes on to the ocean. The latter thus be- 
comes a vast reservoir of water containing salt and other 
mineral matter obtained from the earth's crust. In the 
vital processes of animals and plants it helps change the 
food into a condition fit for distribution and assimilation. 

Industrial Applications. Besides the universal use of 
water for drinking, it is applied to an endless variety of use- 
ful and convenient purposes. It has always been man's 
beast of burden. It is the vehicle for transferring mechan- 
ical energy to water wheels an application now being 
made on a vast scale for generating electricity. It utilizes 
by its peculiar properties the energy in fuel by means of 
the steam engine. It is the highway for transportation on 
the largest scale by ocean, river, lake, and canal. It is the 
vehicle for the distribution of heat by hot water and steam. 
It is the indispensable solvent in metallurgy, in the manu- 
facture of chemicals, and in such industries as soap 
making, bleaching, brewing, dyeing, and tanning; it is 
necessary wherever mortar and cement are used. Man's 
work would be stopped in a thousand other ways were 
he deprived of water. 

Physical Properties of Pure Water. Owing to its 
remarkable solvent power, water is never found pure in 
nature, and is purified even in the laboratory only by taking 
especial precautions. At the ordinary temperature water 
is a tasteless and odorless liquid. It is usually colorless, 
but thick layers are bluish. Water is a poor conductor of 

This last property may be shown by boiling water near the surface 
in a large test tube containing a piece of ice weighted down upon the 
bottom. The ice remains unmelted for some time, although the water 
is boiling a few inches above it. 

34 Descriptive Chemistry. 

Most liquids expand with heat and contract with cold. 
Water is an exception. If water at 100 C. is gradually 
cooled, it contracts in volume. But when 4 C. is reached, 
if the cooling continues, the volume increases as long as 
the liquid state is maintained. Hence at 4 C. a given 
volume contains the greatest weight of water. That is, 
water has its maximum density at 4 C. 

The density of water at 4 C. is i ; and water at this temperature is 
the standard for determining the densities of solids and liquids. Thus, 
when we say the density of gold is 19, we mean that gold is 19 times 
heavier than an equal volume of water at 4C. 

The expansion of water when cooled from 4 C. to o C. is slight, but 
the change is exceedingly important in nature. When the water on the 
surface of a lake or river cools, it contracts, and since it is heavier 
(volume for volume) than the warmer water beneath, it sinks. The 
warmer water rises, is cooled, and likewise sinks, thus causing a circula- 
tion which continues until all the water from surface to bottom has the 
temperature of 4C. Now if the cooling continues, the surface water 
expands and remains on top, because it is lighter than the water 
beneath. Hence when the temperature of the air falls to oC, this top 
layer of water freezes and protects the remaining water from the cold, 
thus stopping the circulation. Should the circulation continue, as the 
temperature fell from 4 C. to o C., the whole body of water would 
finally freeze from top to bottom. This condition would not only 
destroy the fish and marine plants, but seriously affect climate, since 
the heat of summer could not melt such a vast mass of ice. 

When water freezes, it expands about one tenth of its 
volume. That is, 100 cc. of water produce about no cc. 
of ice. In other words, 100 cc. of water and 1 10 cc. of ice 
weigh 100 gm. Hence ice floats. The specific gravity of 
ice is about 0.92. 

The pressure exerted by water when it freezes is powerful. Vessels 
or pipes completely filled with water often burst when the water freezes. 
It is an erroneous but popular idea that " thawing out " a pipe bursts it. 
As a matter of fact, ice contracts when it melts. The pipe cracks when 
the water freezes, and as the ice melts a channel is left for the water to 

Properties of Water. 35 

flow out of the pipe. Because of this property, ice is an effective agent 
in splitting rocks. Water creeps into cracks, especially into the narrow 
ones by capillary attraction, and when it freezes, the rock is slowly split 
apart. Water in freezing also destroys the tissue of living plants, which 
are often said to have been "touched by frost." Frozen flesh for a 
similar reason becomes pulpy and is more liable to putrefy when thawed 
a fact sometimes overlooked by those who eat flesh food which has 
been kept in cold storage. 

FIG. 5. Snow crystals. 
From photographs by Wilson A. Bentley. 

Ice melts at o C. (32 F.), which is also the freezing 
point of water. Ice often crystallizes in freezing, but the 

36 Descriptive Chemistry. 

individual crystals are seldom visible except during the first 
stages of the process. Snow crystals are common (Fig. 5). 
They are always six-sided, and are formed in the atmos- 
phere by the freezing of water vapor. 

Water evaporates at all temperatures, passing off as an 
invisible vapor into the atmosphere or into the air confined 
over it. If water is heated, the vapor passes off rapidly 
until the thermometer reads iooC. (or 212 F.). At this 
point water boils, i.e. it changes rapidly into vapor without 
rise of temperature. This vapor, if allowed to escape into 
the atmosphere, cools and condenses quickly into a cloud 
of minute drops of water. This cloud is popularly called 
steam. Scientifically, steam is invisible. What we call 
steam is a mass of very small particles of water. This 
may be illustrated by boiling water in a large flask. The 
inside of the flask is perfectly transparent, although there 
is a cloud of " steam " issuing from its mouth. 

Water boils when its vapor escapes with sufficient pressure to over- 
come the pressure of the atmosphere upon its surface. Hence the boil- 
ing point depends upon the pressure either of the atmosphere or of 
the vapor within the vessel. The boiling point is iooC. (or 2I2F.) 
when the atmospheric pressure is normal, i.e. 760 mm. The boiling 
point is lower as the pressure is decreased and higher as the pressure is 
increased. Warm water will boil under the receiver of an air pump or 
on the top of a high mountain. In the city of Mexico (7500 feet above 
sea level) water boils at about 92 C., and in Quito in South America 
(9350 feet above sea level) water, which boils at about 90 C., is not 
hot enough to cook potatoes. 

The pressure which water vapor exerts as it escapes from a liquid is 
called its vapor tension. Since the rate of evaporation depends upon 
the temperature of the liquid, vapor tension varies with the temperature. 
Vapor tension is usually expressed in millimeters of mercury. Thus, at 
iooC. the vapor tension of water is 760 mm., because at the boiling 
point the vapor pressure is just enough to overcome the opposing 
atmospheric pressure. At 20 C. the vapor tension of water is 17.39 mm - 

Properties of Water. 37 

A liter of steajn, if it could exist at oC. and 760 mm. pressure, would 
weigh 0.806 gm., or nine times more than a liter of hydrogen. 

Natural Waters. Water is never found pure in nature. 
Even rain water, which is usually regarded as the purest 
natural water, contains gases and dust washed from the 
air. When rain strikes the ground it begins at once to 
take up impurities from the rocks, soil, and vegetation. 
Some of the water flows along the surface, becoming more 
and more impure, and finally reaches the ocean. From 25 
to 40 per cent of the annual rainfall in temperate regions 
soaks into the ground and trickles through the soil at an 
estimated rate of 0.2 to 20 feet a day. This underground 
water finally finds its way again to the surface as a spring 
or well, through a lake or river, or from a hillside. On its 
journey underground the water loses most, often all, of its 
organic matter, remnants of vegetable and animal matter, 
but dissolves mineral matter and gases. If the amount 
of dissolved matter in spring water is large or the kind of 
matter is so unusual as to give the water a marked taste or 
medicinal properties, the water is called mineral water. 
Water containing calcium and magnesium compounds is 
hard, but in soft water, such as rain water, these com- 
pounds are absent. 

There are several hundred mineral springs in the United States. 
Those having a high temperature are called thermal, as at Hot Springs, 
Arkansas, and at Bath, England. Many contain a large proportion of 
common salt, as at Saratoga, New York. Others contain alkaline matter 
and carbon dioxide gas, eg. Vichy and Apollinaris water. Sulphur 
springs contain solid or gaseous compounds of sulphur or both and 
have valuable medicinal properties. Some, like Hunyadi, are bitter; 
but others, especially those in New York State, which contain gaseous 
sulphur compounds, have a sweet taste but an unpleasant odor. Cha- 
lybeate waters contain soluble iron compounds. Many waters contain 
lime and magnesium compounds, and a few contain alum. Most natural 

Descriptive Chemistry. 

mineral waters contain traces of a large number of different substances. 
Many commercial mineral waters have doubtful medicinal value. 

River water obviously contains the impurities brought 
by springs and the surface water ; it is also often made 
very impure by decaying animal and vegetable matter, 
which has been purposely or accidentally introduced, espe- 
cially if the river passes through a thickly settled region. 
A sluggish river is more apt to be impure than a swift 
one, because the latter tends to purify itself by exposing 
its impurities to the oxidizing power of the air. Ocean 
water contains a large proportion of common salt. The 
other substances in order of abundance are magnesium 
chloride, magnesium sulphate, calcium sulphate, and potas- 
sium chloride ; many other substances are present in small 
quantities. The peculiar taste of ocean water is due to 
the presence of these substances, and since by evaporation 
the water only is removed, the ocean always has a " salty " 
taste. The composition of some natural waters is sum- 
marized in the following 
















Rain . 





6. 4 











Spring . . 




I 5 .8 



Mineral (Bath) 









Ocean . . 






12. 1 



Properties of Water. 39 

Drinking Water. " Water used for drinking should be free from 
visible suspended particles, without disagreeable taste or smell, and not 
capable of acquiring such by standing for a day or two in a clean, well- 
closed vessel. It should also contain enough of the gases derived from 
the atmosphere to give it a slight taste distinguishable from the flatness 
of boiled or distilled water. It should not contain solid matter in solu- 
tion to the extent of more than 300 parts in a million, or about 3 gm. 
to 10 1. The mineral portion of this solid matter should not con- 
tain any poisonous substance. As little as possible of the solid contents 
should consist of organic matter usually not exceeding 15 to 20 
parts per million, or about 2 gm. to 100 1. And it is particularly de- 
sirable that decomposing animal matter or substances which give evi- 
dence of its previous presence should be found, if at all, as a mere trace. 
Above all, drinking water should be free from disease-producing bacte- 
ria or other injurious microorganisms." 

The problem of obtaining drinking water in large quantities is usually 
local. In some cities the water is purified by filtering it through a 
layer of sand and gravel, an acre or more in area and several feet deep. 
Such a filter removes bacteria almost completely, though it must be 
frequently cleaned. Sometimes the water is stored in a large settling 
basin or reservoir and purified by adding alum, or a similar substance, 
which causes the suspended matter to settle. Dissolved substances 
cannot be removed without considerable difficulty, so as a rule water 
is taken from a source which is reasonably pure. 

The purity of drinking water is usually determined by a water analy- 
sis. This is not a decomposition of water, but a chemical examination 
of a sample for the presence and amount of certain substances which 
indicate or cause impurity. A chemical examination is of limited value, 
however, unless it is supplemented by a microscopic study of a fresh 
sample and a rigid sanitary inspection of the premises. Water which 
is clear, sparkling, cool, attractive to the eye, and pleasant to the taste 
may be seriously polluted by disease germs, or may be liable to sudden 
contamination from some unsuspected source. On the other hand, a 
rather unpleasant-looking water may be harmless. Hence the necessity 
of careful and extended examination of water to be used as a beverage. 

Water may be purified by distillation. This operation is not con- 
venient with large quantities. It is performed in the laboratory in a 
condenser, which is shown in Figure 6 arranged for use. 

The condenser consists of an outer tube, A A, provided with an inlet 

4 o 

Descriptive Chemistry. 

and an outlet for a current of cold water, which surrounds the inner 
tube, BB. The vapor from the water boiling in the flask, C, condenses 

FlG. 6. Condenser arranged for the distillation of water. 

in the inner tube, owing to the decrease in temperature, and drops off 
the lower end of this tube, as the distillate, into the receiver, D, while 
the impurities remain behind in the flask. Distilled water is prepared 
on a large scale in metal vessels, and the vapor is con- 
densed in a block tin pipe coiled around the inside of 
a vessel through which a 
current of cold water is flow- 
ing. This coiled pipe is 
called a worm (Fig. 7). 
Distilled water is used in 
the chemical laboratory ; 
large quantities are made 
into ice. Distillation is an 
old process. A quaint still 
is shown in Figure 8. Dis- 
tillation is the process used 
to separate liquids from 
solids and from each other, 
FIG. 7. Worm- and finds extensive appli- FIG. 8. A quaint still, 
shaped tube. cation in the manufacture of liquors and kerosene oil. 

Properties of Water. 41 

Solution. Many solids, liquids, and gases disappear 
when put into water. This operation is called dissolving, 
or putting into solution. The resulting liquid is called a 
solution of the substance used. The liquid in which the 
substance dissolves is called the solvent, and the dissolved 
substance is called the solute. If the solute is not vola- 
tile, or not very volatile, it may be recovered by evaporat- 
ing, or distilling off, the water. The degree of solubility 
is usually expressed by the terms sligJitly soluble, soluble, 
and very soluble. It is more accurate, and usually desir- 
able, to state the proportions of solvent and solute, and 
also the temperature. Thus, instead of saying that 
common salt is very soluble in cold water, it is better to 
state that 36 gm. of salt dissolve in 100 cc. of water at 
20 C. Substances which do not dissolve in water are 
called insoluble, though this term is also applied to those 
substances a minute quantity of which dissolves in water. 
Thus glass, sand, and many rocks are usually classed as 
insoluble substances, but they dissolve appreciably in 

A solution which contains a small proportion of solute 
is called a dilute solution ; one containing a large propor- 
tion is called a concentrated solution. Thus, dilute sul- 
phuric acid usually contains one volume of acid to three 
or more volumes of water, while concentrated sulphuric 
acid is nearly 98 per cent acid. Sometimes the terms 
weak and strong replace dilute and concentrated, but they 
are ambiguous, and their use should be avoided. 

Solutions of Gases. Water dissolves or absorbs many 
gases. The degree of solubility depends upon the gas, 
the temperature of the water, and the pressure at which 
solution occurs. Some gases, such as ammonia and hydro- 
chloric acid gas, are very soluble in water. Advantage of 

Descriptive Chemistry. 

this fact is taken in manufacturing ammonium hydroxide 
and hydrochloric acid. Each commercial substance is 
merely a water solution of the respective gases, ammonia 
and hydrochloric acid gas ; the gas is readily liberated by 
heating the liquid. 

The common gases, oxygen and hydrogen, are only slightly soluble 
in water. Air dissolves in water, as may easily be shown by heating 
faucet water, bubbles of air forming and escaping quickly as heat is ap- 
plied. Carbon dioxide gas is very soluble in water. Water containing 
this gas is called "soda water," or carbonated water. More gas is forced 
into the water than will dissolve at the ordinary temperature and pres- 
sure, as may be seen by the rapid escape of gas when the water is drawn 
from a soda fountain. This rapid escape of a gas is called efferves- 
cence. " Soda water " must, therefore, be stored in a strong vessel and 
kept in a cool place. The gas was formerly obtained from sodium bi- 
carbonate a compound related to "soda' 1 ; hence the name "soda 
water. 11 It is now prepared from marble and an acid, or from liquid 
carbon dioxide. 

The volume of gas which will dissolve in water decreases with rise 
of temperature. Thus, 100 cc. of water at oC. will dissolve 179.6 cc. 
of carbon dioxide, but only 90.1 cc. at 20 C. The volume of a mod- 
erately soluble gas which is dissolved by water is directly proportional 
to the pressure if the temperature is constant. This is Henry's law. 
It is illustrated by the following 





I 1. 

900 cc. 
1800 cc. 
3600 cc. 



7200 cc. 


The tremendous pressure to which subterranean gases are subjected 
accounts for their presence, especially carbon dioxide, in such large pro- 
portions in the waters of mineral springs. 

Properties of Water. 43 

Solutions of Liquids. The solubility % of liquids in 
water varies between wide limits. Some, such as alcohol 
and glycerine, are soluble in all proportions. Oils, such 
as kerosene, are practically insoluble; hence the old adage, 
" Oil and water will not mix." Carbon disulphide is also 
insoluble, as may be seen by the formation, after agitation, 
of two distinct layers of liquid. The existence of two 
layers, however, is not always absolute proof of insolubility. 
JEther and water form two layers, but each dissolves appre- 
ciably in the other. In many cases a rise of temperature 
increases the solubility of liquids in water. 

Solutions of Solids. The solubility of solids in water 
is a matter of tremendous practical importance. The 
abundance of water and its power to dissolve such a vast 
number of different solids have led some to call water 
" the universal solvent." The far-reaching effect of this 
marvelous power in nature and its indispensable value to 
man have been considered. (See above.) 

The degree of solubility of solids in water varies with 
the substance and with the temperature of the water. 
Some, like potassium permanganate, are very soluble, while 
others, like calcium sulphate, are difficultly soluble. In 
most cases solubility increases with a rise of temperature ; 
hence the common practice of heating to hasten solution. 
The effect of increased temperature on solubility is some- 
times very marked, the solubility being increased fourfold 
in some cases. Calcium hydroxide is less soluble in hot 
than in cold water, while common salt (sodium chloride) 
dissolves to about the same degree in each. There is a 
limit to solubility. That is, a given volume of water at 
a fixed temperature will dissolve a definite weight of solid 
and no more, although some undissolved solid remains in 
the water. 

44 Descriptive Chemistry. 





20 C. 


Calcium chloride 



Copper sulphate (cryst.) . . 



Magnesium sulphate . . 



Potassium chlorate .... 



Potassium chloride .... 



Potassium dichromate . . . 



Potassium nitrate . . . 


2 4 6 

Potassium sulphate .... 



Sodium chloride 



A solution is saturated at a given temperature when it 
will dissolve no more solid. If a hot solution, especially 
one which contains much solid, is cooled slowly, the solid 
soon begins to separate from the liquid, since solubility 
usually decreases with a fall of temperature. Often the 
solid is deposited in masses having a definite shape. This 
operation is called crystallization, and the masses are 
called crystals (see below). The shape and color of the 
crystal are characteristic of the substance, and serve to 
identify it. Thus, common salt crystallizes in cubes. 
Sometimes it is more convenient to evaporate a hot, con- 
centrated solution. The point of saturation at the lower 
temperature is thus reached so gradually that the crystals 
can grow symmetrically. A brief account of crystals will 
be found in 3 of the Appendix. 

A solid can also be separated from a solution by precipi- 
tation. This may be done in two ways, (i) By adding 
a liquid in which the solid is not very soluble. Thus, when 

Properties of Water. 45 

water is added to an alcoholic solution of camphor, the 
liquid becomes turbid, or cloudy, because the camphor is 
not soluble in water. That is, the solid has been precipi- 
tated as very fine particles which remain suspended in the 
liquid for some time. Since the separated solid sooner or 
later falls to the bottom of the vessel, it is called a precipi- 
tate. (2) By changing the dissolved solid into another 
substance not soluble in the liquid. Such chemical changes 
are examples of double decomposition. Thus, when so- 
dium chloride solution is added to silver nitrate solution a 
white, curdy precipitate of silver chloride is formed. A 
soluble silver compound has thus been changed into an 
insoluble silver compound, thereby removing the combined 
silver from the solution. So, also, a soluble chlorine com- 
pound (sodium chloride) has been changed into an insoluble 
chlorine compound (silver chloride), thereby removing the 
combined chlorine from the solution. Precipitation is a 
very common operation in chemistry. 

A hot, saturated solution of some solids, such as sodium 
sulphate and sodium thiosulphate, deposits no crystals 
when the clear solution cools. Such solutions are super- 
saturated. Supersaturation can occur only when the un- 
dissolved solid is not present. Hence, if a fragment of 
the solid is dropped into the supersaturated solution, crys- 
tals soon begin to form upon the fragment, and this sepa- 
ration continues until nearly all the substance is deposited, 
often forming a solid mass in the test tube. Dust, or even 
shaking, causes the substance to be deposited, hence the 
solution should be kept corked and left undisturbed. Sat- 
uration is analogous to stable equilibrium, while supersatu- 
ration resembles unstable equilibrium. 

Water of Crystallization. Crystals deposited from 
the water solution of many solids, even after they are dried 

4.6 Descriptive Chemistry. 

by pressing between filter paper or by exposure to a mod- 
erate temperature, often contain water which seems to be 
an essential part of the chemical compound. This water 
is called water of crystallization. The crystals of some 
compounds, e.g. sodium carbonate and sodium sulphate, 
lose their water of crystallization and crumble on exposure 
to the air. This property is called efflorescence, and such 
crystals are said to effloresce or to be efflorescent. Heat 
will drive the water of crystallization from crystals which 
contain it, e.g. gypsum, alum, and copper sulphate. 

The proportion of water of crystallization in crystals is not arbitrary. 
It is constant in the same compound when crystallized under uniform 
conditions, but the proportion varies between wide limits in different 
substances. No explanation has been given of the varying amount of 
water of crystallization, nor of its necessity for the form and color of 
some crystals and not for others. Some well-crystallized substances 
contain no water of crystallization, e.g. potassium nitrate, potassium 
dichromate, sugar, and salt. 

Crystals which have lost their water of crystallization are said to be 
dehydrated or anhydrous. Thus, the grayish powder obtained by 
heating the blue crystallized copper sulphate is called dehydrated cop- 
per sulphate. The words dehydrated and anhydrous have been extended 
to describe any substance from which water has been removed, as anhy- 
drous alcohol or ether. The opposite term, hydrated, is sometimes 
applied to a compound to emphasize the fact that it contains water of 

Deliquescence. Many substances, crystallized and 
uncrystallized, absorb water when exposed to the air, and 
become moist, or even dissolve in the water. Calcium 
chloride, potassium carbonate, zinc chloride, sodium hydrox- 
ide, and potassium hydroxide belong to this class. This 
property is called deliquescence, and the substances are 
said to deliquesce, or to be deliquescent. The term hygro- 
scopic is applied to substances which absorb water, but 
hygroscopic substances do not dissolve in the absorbed 

Properties of Water. 47 

water, and sometimes do not even become moist. Quick- 
lime is hygroscopic. 

Common salt, or sodium chloride, often appears to deliquesce, espe- 
cially in damp weather. The deliquescence is due, however, to the 
presence of magnesium and calcium chlorides. Sodium nitrate is some- 
what deliquescent, and cannot be used in the manufacture of gunpowder, 
so potassium nitrate is used instead. This property of deliquescence is 
often utilized in the laboratory to remove water vapor from gases, cal- 
cium chloride being especially serviceable for this purpose. 

Thermal Phenomena of Solution. Solution is often accompanied 
by an appreciable change of temperature. When sulphuric acid is poured 
into water, heat is produced. With large quantities the heat is so great 
that the mixture often boils, and sometimes the hot acid is spattered. 
Hence, the acid should be added slowly to the water, and the mixture 
constantly stirred. Other substances which dissolve with the liberation 
of heat are fused calcium chloride, potassium hydroxide, and sodium 
hydroxide. Some which dissolve with a fall of temperature are crystal- 
lized calcium chloride, ammonium nitrate, ammonium chloride, and 
potassium nitrate. This subject is still under investigation. 

Solution and Chemical Action. Probably when a sub- 
stance dissolves it is so modified that it can participate 
more readily in chemical changes. Hence, solution is an 
aid to chemical change, and is often an easy means of 
causing it. Thus, if dry tartaric acid and sodium bicar- 
bonate are mixed, there is no evidence of chemical action ; 
but when the mixture is poured into water, the copious 
evolution of carbon dioxide gas is conclusive evidence of a 
chemical change. Similarly, when a dry mixture of ferrous 
sulphate and potassium ferrocyanide is poured into water, 
the immediate appearance of a blue precipitate shows that 
the water was needed for the chemical change. Solution 
is such an important aid to chemical action that many 
substances employed in the laboratory are in solution, 
and many processes in chemistry are " wet " processes. 

48 Descriptive Chemistry. 

Mention has already been made of the application of this 
fact to many industries. 

The Nature of Solution has long been a subject of specu- 
lation and study. The problem as a whole is still unsolved, 
though much light has been thrown upon the question by 
recent investigations (see Chapter X). 


1 . Mention several familiar properties of water. 

2. In what forms does water exist ? 

3. Give the per cent of water in some familiar foods. 

4. Develop the topics : () water is an erosive agent ; () water 
is a solvent in nature ; (V) water has many industrial applications ; 

(d) water behaves exceptionally when heated from o C. to ioC.; 

(e) ice floats ; (/) water is a cleansing agent. 

5. Explain these expressions : (a) water has its maximum density 
at 4 C. ; (b) the density of ice is 0.92 ; (V) steam is invisible ; (d) the 
lower the pressure, the lower the boiling point ; (e) 10 cc. per liter ; 
(_/") parts per million. 

6. How do natural waters illustrate the solvent power of water? 

7. What is (a) mineral water, () soft water, (c) hard water, 
(d) sulphur water, (e} chalybeate water ? 

8. What does ocean water contain? Why is the sea water salt? 

9. What constitutes a safe drinking water? How may city water 
be purified ? What is a water analysis ? 

10. Describe the operation of distillation. What is a condenser 
and why is it" so named? Is distillation a new or an old process? 
Of what industrial use is it? 

11. Define and illustrate (a) water of crystallization, (b) efflores- 
cence, (c} deliquescence, (d) hygroscopic, (e) anhydrous, (/) dehy- 
drated, (g) crystal, (//) crystallization. 

12. Define and illustrate (a} solution, (b} solvent, (V) solute, 
(d) soluble, (e) slightly soluble, (/) very soluble, (g} insoluble, (h} di- 
lute, (/) concentrated, (/) saturated solution, () supersaturated 

13. Give several facts about the solubility of gases in water. What 
is (a) soda water, (b} carbonated water? How do we know that air 

Properties of Water. 49 

dissolves in water? Why do subterranean waters often contain dis- 
solved gases? State Henry's law of the solubility of gases. What 
effect has (a) heat and () cold on the solubility of gases in water? 

14. What liquids are soluble in water? How may such liquids be 
separated from water? 

15. What general effect has (a} heat and () cold on the solubility 
of solids in water? Mention some solids which are (a) very soluble, 
(b) moderately soluble, (<:) almost insoluble in water. Develop the 
topic : There is a limit to the solubility of solids in water. 

1 6. (a} How would you find the approximate amount of water in 
(i) milk, (2) an apple? (b) How would you find the per cent of each 
substance in a mixture of sand and sugar? 

17. Develop the topic: Solution aids chemical change. Why are 
so many solutions used in a laboratory ? 

1 8. What changes in volume occur when (a} ice melts, () water 
freezes, (c) water is heated from oC. to i5C., (d) water is cooled 
from I5C. to oC.? 

19. Write an essay on "Mineral Springs in the United States." 


1. If 1.5 gm. of crystallized barium chloride lose 0.22 gm. when 
heated to constant weight, what per cent of water of crystallization 
does it contain? 

2. If 2 gm. of another lot of barium chloride lose 0.295 gm., 
what per cent of it was water of crystallization ? 

3. If a liter of sea water has a density of 1.25, how many grams of 
"salt" does it contain? 

4. If the density of ice is 0.92, what volume will a liter of water at 
4C. occupy when frozen? Ans. 1.087 ! 

5. How much water (approximately) is contained in (a) 2 Ib. of 
lobster, (b} 56 Ib. of potatoes, (c) i Ib. of tomatoes, {d} 2 Ib. of milk, 
(e) i Ib. of white bread, (/) a human body weighing 150 Ib. ? 

6. If a kilogram of sea water contains 36.4 gm. of "salt," what 
per cent of the water is " salt " ? 

7. If a block of ice weighs 280 kg., what is its volume? 

Ans. 304.3 1. 

8. A solution measures 100 cc. and contains 15 gm. of potassium 
nitrate. What per cent of water and of solid is in the solution ? 


- WATER was considered an element until about the end 
of the eighteenth century. At that time it was shown to 
be a compound of hydrogen and oxygen. Many famous 
chemists worked on this problem. 

The Composition of a Compound is determined either 
by analysis or synthesis, i.e. by taking it apart or putting 
its parts together. Sometimes both methods are used, 
since each method fortifies the other and strengthens the 
final conclusion. These methods find excellent application 
in determining the composition of water. 

Analysis and synthesis may be qualitative or quantitative. A quali- 
tative experiment is a study of the properties of elements and com- 
pounds with a view of discovering what they contain. A quantitative 
experiment is an accurate determination of the weight or volume of the 
components of a compound. Qualitative tests involve merely quality, 
while in quantitative tests quantity is the essential feature. Obviously, 
a complete determination of the composition of a compound requires 
both tests. 

Water contains Hydrogen. When steam is passed 
over heated metals, hydrogen is liberated. Lavoisier's 
demonstration of this fact has already been considered 
(see Preparation of Hydrogen). The fact that sodium 
liberates hydrogen from water at the ordinary temperature 
has also been discussed (see ibid.). If red litmus paper is 
put into the water from which the sodium has liberated 
hydrogen, the litmus paper becomes blue. This change 


Composition of Water. 51 


of color from red to blue shows that an alkali is in the 
water, because alkalies turn red litmus paper blue. The 
alkali is sodium hydroxide, and it may be obtained as a 
white solid by evaporating the water. Sodium hydroxide 
is a compound of sodium, hydrogen, and oxygen, and is 
formed by replacing part of the hydrogen of water by 
sodium. Since sodium liberates hydrogen from water, and 
forms at the same time a compound sodium hydroxide 
containing hydrogen, the hydrogen in water must be 
divisible into two parts. Now if o. I gm. of sodium is 
allowed to act upon water, 48.22 cc. of hydrogen are liber- 
ated ; and if the sodiunThydroxide thus formed is dried and 
heated with sodium, 48.22 cc. more of hydrogen are ob- 
tained. This shows that the hydrogen in water is divisible 
into two equal parts a fact which will soon be utilized. 

Water contains Oxygen. The fact that oxygen is a 
component of water has already been suggested, e.g. (i) 
by the production of water when hydrogen is burned in 
air, (2) by the formation of a compound of iron and oxy- 
gen when steam is passed over hot iron, and (3) by the 
formation of sodium hydroxide when sodium acts upon 
water. These proofs, however, are all indirect. A simple 
direct demonstration of the presence of oxygen in water 
may be made by allowing chlorine water to stand in the 
sunlight. (Chlorine water is prepared by saturating water 
with chlorine gas an element to be studied in Chapter 

XI.) A long tube like that shown ^ 

in Figure 9 is completely filled GU 

with chlorine water, the open end is FIG. 9. Tube for decompo- 

, L . . sition of water by chlorine. 

immersed in a vessel containing 

some of the* same solution, and the whole apparatus is 
placed in the direct sunlight. Bubbles of gas soon appear 
in the liquid, and after a few hours a small volume of 

Descriptive Chemistry. 

gas collects at the top of the tube. This gas may be 
shown, by the usual tests, to be oxygen. 

The Electrolysis of Water is its decomposition by elec- 
tricity. It is accomplished in the apparatus shown in 

Figure i o. Since pure water does 
not conduct electricity, sulphuric 
acid is added. Enough of this 
acid mixture is poured into the 
apparatus to fill the reservoir 
half full after the stopcocks have 
been closed. As soon as an 
electric battery of two or more 
cells is connected by wires with 
the piece of platinum near the 
bottom of each tube, bubbles 
of gas form on the platinum, 
and as the action proceeds, the 
bubbles rise and displace the 
water in each tube. The volume 
of gas is greater in one tube. 
Assuming that the tubes have 
the same diameter, the volumes 
are in the same ratio as their 
heights, which will be found by 
measurement to be two to one. 
The larger volume of gas is 
FIG. io. Hofmann apparatus for hydrogen and the smaller one 

electrolysis of water. 

is oxygen. Many accurate repe- 
titions of this experiment have shown that only hydrogen 
and oxygen are produced, and that the ratio of their volumes 
is two to one. It has also been shown that the sum of the 
weights of the two gases equals the weight of the water 
decomposed. The whole experiment demonstrates that 

Composition of Water. 53 

water is a compound consisting of two volumes of hydro- 
gen combined with one volume of oxygen. 

Water was first decomposed by electricity in 1800 by Nicholson and 
Carlisle. Davy confirmed their work by a series of brilliant experi- 
ments extending through a period of six years (1800-1806). During 
this time he not only proved that the volume of hydrogen is double that 
of oxygen, but by electrolyzing water in a gold vessel placed in an atmos- 
phere of hydrogen, he proved that nothing but these gases is produced. 

The Quantitative Composition of Water. The fore- 
going facts about the composition of water have been 
mainly qualitative. They have shown by analysis and 
synthesis that water consists of hydrogen and oxygen, and 
that the ratio of their volumes is approximately two to one. 
Decisive evidence of the quantitative composition of water 
is obtained by a determination of its volumetric and its 
gravimetric composition. Volumetric means "by volume" 
and gravimetric means " by weight." 

The Volumetric Composition of Water is determined 
by exploding a mixture of known volumes of hydrogen 
and oxygen in a eudiometer. 

Gas volumes which are to be compared with each other must be dry 
and at the same temperature and pressure. This requirement, which is 
called the " standard condition," is inconvenient, and almost impracti- 
cable. Hence, it is customary to measure each volume of moist gas 
under the existing conditions, and then reduce the observed volume to 
that volume which the gas would occupy if standard conditions pre- 
vailed. The reduction to standard conditions is accomplished by the 
formula j/r /pi _ n \ 

760(1 + . 00366 /) 

In the formula l V = the corrected volume. 
V = the observed volume. 

1 A complete discussion of the laws of gases, the principles which control their 
measurement, together with the development of the above formula for reduction to 
standard conditions, may be found in Appendix B of the author's " Experimental 
Chemistry." See also the Laws of Boyle and Charles in Chapter II, and Vapor 
Density in Chapter IV (this book). 


Descriptive Chemistry. 

P f = the observed pressure. 
/ = the observed temperature. 
a the vapor tension at / C. 

A convenient form of apparatus for determining the volu- 
metric composition of water is shown in Figure n. The 
essential part is the eudiometer, F. In this graduated 
glass tube the gases are accurately measured and ex- 
ploded. The electric spark which causes the explosion is 

obtained from an induc- 
tion coil and battery. 
The spark leaps across 
the space between the 
platinum wires at the 
top of the eudiometer, 
and the heat produced 
by this spark causes the 
hydrogen and oxygen 
to combine and form 
water. Oxygen and hy- 
drogen are introduced 
separately into the eudi- 
ometer, measured, and 
- exploded. After the 

FIG. ii. Apparatus for determining the volu- explosion, which IS indi- 

cated by a slight click 

or flash of light, water from the reservoir, E, rushes up 
into the eudiometer. The water does not completely fill 
the tube, because an excess of one gas is added. This 
additional gas takes no part in the chemical change, but 
merely serves to lessen the violence of the explosion, which 
otherwise might break the eudiometer. The quantity of 
water formed by the union of the hydrogen and oxygen 

Composition of Water. 


is too minute to measure. Repeated trials of this experi- 
ment show that two volumes of hydrogen always combine 
with one volume of oxygen. This is the volumetric com- 
position of water. 

The discovery of the volumetric composition of water was not 

made by one chemist alone. Priestley, about 1780, noticed that when 
a mixture of air and hydrogen was exploded, " the inside of the glass, 
though clear and dry before, immediately became dewy." Cavendish, 
in 1781, showed that when a mixture of two parts hydrogen and one 
part oxygen was exploded, nothing but water was formed. Watt, in 
1783, was the first to state that water is a compound, though he per- 
formed no experiments and probably did not understand the real nature 
of its components. Lavoisier in the same year verified many facts pre- 
viously noticed but not completely understood, and undoubtedly first 
clearly recognized and stated what his contemporaries had overlooked. 
The final proof of the volumetric composition of water was an accurate 
verification in 1805 by Gay-Lussac and Humboldt of the previous ob- 
servation that two volumes of hydrogen unite with one volume of oxygen. 

The Gravimetric Composition of Water is determined 
by passing dry hydrogen over copper oxide. The method 
depends upon the fact that many oxides, such as those of 
lead, copper, and iron, when heated in a current of hydro- 

=7) (e= = l 



c c' 


FIG. 12. Apparatus for determining the gravimetric composition of water. 

gen, give up their oxygen, or, chemically speaking, these 
oxides are reduced to metals. By this reduction the oxy- 
gen of the oxide combines with the hydrogen, thereby 
forming water which is collected in a weighed tube. 

56 Descriptive Chemistry. 

A convenient form of apparatus is shown in Figure 12. 
The copper oxide is placed in the combustion tube, CC, 
which is made of hard glass. The Marchand tube, D, 
which is filled with calcium chloride, collects and retains 
the water formed in the combustion tube, as the hydrogen 
passes over the hot copper oxide. The tubes A, B, and E 
keep moisture out of the apparatus. The experiment is 
very simple. Copper oxide is placed in the combustion 
tube, which is then carefully weighed. The Marchand 
tube, being filled with calcium chloride, is also weighed. 
After the other tubes are properly filled and the hydrogen 
generator adjusted, the tubes are connected as shown in 
the figure. The combustion tube is now heated, and mois- 
ture collects in it; as the heat increases the copper oxide 
glows, and the moisture passes into the Marchand tube. 
When the operation is over and the apparatus is cool and 
free from hydrogen, the combustion tube and Marchand 
tube are weighed. The gain in weight of the Marchand 
tube is the weight of the water formed, while the loss in 
weight of the combustion tube is the weight of the oxygen 
contained in this water. An illustration will make this 
clear. Dumas and Stas, who performed this experiment 
accurately in 1843, found substantially that the combus- 
tion tube lost 5.251 gm. of oxygen, while the Marchand 
tube gained 5.909 gm. of water. But 5.251 and 5.909 
are in the same ratio as 8 and 9. Thus : 

5.251 : 5.909 : : 8 : 9. 

This means that oxygen makes up f of water. The re- 
maining ^ is of course hydrogen. In other words, the 
gravimetric composition of water is eight parts oxygen 
and one part hydrogen. This ratio is often stated in per- 
centage ; thus water contains 

Composition of Water. 57 

88.88 per cent of oxygen. 
1 1 . 1 1 per cent of hydrogen. 

For reasons which will soon be given, it is more conven- 
ient to state the composition of water by weight, as two 
parts hydrogen to sixteen parts oxygen. 

The gravimetric composition of water was first determined about 
1820 by Berzelius and Dulong. Their work was verified by Dumas 
and Stas in 1843. 

A Comparison of the Volumetric and Gravimetric Com- 
position of Water shows that the results of the two 
methods agree. One volume of oxygen is sixteen times 
heavier than an equal volume of hydrogen (see Density of 
Hydrogen). Therefore, the one volume of oxygen must 
be eight times heavier than the two volumes of hydrogen 
in water. That is, the oxygen in water weighs eight times 
more than the hydrogen. But this is the ratio actually 
found in determining the gravimetric composition of water 
by an independent experiment. These facts strengthen 
our belief that the composition of water is 

By weight, one part hydrogen and eight parts oxygen. 
By volume, two parts hydrogen and one part oxygen. 

Summary. The following facts have been shown con- 
cerning the composition of water : 

(1) Water is a chemical compound 'of hydrogen and 

(2) It is formed when hydrogen is burned in air, or 
when a mixture of hydrogen and oxygen is exploded. 

(3) It can be decomposed by electricity into hydrogen 
and oxygen in the ratio of two volumes of hydrogen to one 
volume of oxygen. 

58 Descriptive Chemistry. 

(4) Sodium liberates hydrogen from water and forms at 
the same time a solid containing a quantity of hydrogen 
equal to the quantity of hydrogen liberated. Iron, other 
metals, and carbon liberate hydrogen from water, forming 
at the same time an oxide of the respective substance. 

(5) Chlorine liberates oxygen from water. 

(6) Two volumes of hydrogen, when exploded with one 
volume of oxygen, combine to form water, and the weight 
of the water formed equals the weight of the gases used. 

(7) Water is formed by the union of two parts by weight 
of hydrogen and sixteen parts by weight of oxygen. 


1. How is the composition of a compound determined ? 

2. Define (a) synthesis, () analysis, (<:) qualitative, (//) quantita- 
tive, (e) volumetric, (/) gravimetric. 

3. How would you prove that water is composed of hydrogen and 
oxygen ? 

4. How do we know that the hydrogen in water is divisible into two 
equal parts ? 

5. What is the electrolysis of water ? How is it accomplished ? 
What does it prove about the composition of water ? When and by 
whom was it first performed ? What did Davy contribute toward the 
solution of the problem ? 

6. What is the volumetric composition of water ? How is it deter- 
mined ? Who worked on this problem, and what did each contribute 
to its solution ? 

7. Answer the same questions (as in 6) about the gravimetric com- 
position of water. 

8. Compare the volumetric and the gravimetric composition of 

9. What does the burning of hydrogen show about the composition 
of water ? 

10. Summarize the essential facts regarding the composition of 

Composition of Water. 59 

ii. Give a brief biographical account of (a) Nicholson and Carlisle, 
() Dumas, (c} Humboldt, (d} Stas, (e) Watt, (/) Gay-Lussac (see 
Appendix, 4) . 


1. What weight of (a) hydrogen and (fr) oxygen can be obtained 
by decomposing 125 gm. of water ? 

2. What volume of (a) hydrogen and (6) oxygen can be obtained 
by decomposing 9 1. of water ? 

3. What weight of hydrogen must unite with 16 gm. of oxygen to 
form water ? What weight with (#) 40 gm., (b) 70 gm., (c) 160 gm. ? 

4. What volume of oxygen must unite with 2 1. of hydrogen to form 
water ? What volume with (a) 40 1., () 40 cc., (c) 40 qt, (d ) 95 vol- 
umes, (e) 1 60 1. ? 

5. What volume of oxygen is necessary to unite with 100 gm. of 
hydrogen to form water ? (Suggestion : What is the weight of a liter 
of oxygen ?) 

6. Hydrogen is passed over 2.48 gm. of hot copper oxide, which at 
the end of the experiment weighed 2.24 gm. ; the water formed weighed 
0.27 gm. In what ratio did the hydrogen and oxygen combine ? 

7. Berzelius and Dulong, in 1820, obtained the following results in 
their determinations of the gravimetric composition of water : Loss of 
weight of copper oxide (in grams), 10.832 and 8.246. Weight of water 
formed, 12.197 and 9.27. Calculate in each case the ratio in which the 
hydrogen and oxygen combined. What is the average ratio ? 

8. Dumas and Stas repeated the above work in 1843, and found as 
an average of nineteen determinations, that 840.161 gm. of oxygen 
formed 945.439 gm. of water. Calculate the ratio of combination. 

Hydrogen Dioxide is a liquid composed of hydrogen and oxygen. 
But the proportion of the components is not the same as in water. It 
contains two parts of hydrogen and thirty-two parts of oxygen by 
weight. It is often called, especially in commerce, hydrogen peroxide, 
because its relative proportion of oxygen is greater than in water the 
other hydrogen oxide. 

It is manufactured by treating barium dioxide (or peroxide) with 
sulphuric or hydrochloric acid. The commercial solution has a vari- 
able strength, and usually contains three or more per cent of hydrogen 
dioxide. It has a sharp, pungent odor, and a bitter, metallic taste. 

60 Descriptive Chemistry. 

Hydrogen dioxide is an unstable compound ; it decomposes slowly at 
the ordinary temperature, and very rapidly if heated. The dilute, com- 
mercial solution is somewhat stable, but heat decomposes it completely 
into water and oxygen. The ease with which it yields oxygen makes 
it a good oxidizing agent. In this respect, hydrogen dioxide resembles 
ozone, and, indeed, they are sometimes mistaken for each other. It is 
also a reducing agent, and is frequently used as such in the laboratory. 
It is used extensively to bleach animal and vegetable matter, such as 
human hair, ostrich feathers, fur, silk, wool, cotton, bone, and ivory. It 
is also used as an antiseptic and disinfectant in surgery. Large quanti- 
ties are used to restore the color to faded paintings a use suggested 
by The'nard, the discoverer. In the laboratory it is proving a service- 
able reagent. 

Hydrogen dioxide is found in the air, in rain and snow, but the 
proportion is variable and exceedingly small. 


The Atmosphere is the great mass of gas surrounding 
the earth and extending into space. Its estimated height 
is fifty to several hundred miles. We live at the bottom 
of this vast ocean of air, as it is often called. 

Aristotle (384-322 B.C.) regarded air as one of the four elementary 
principles whose combinations made up all substances in the universe. 
The other three were earth, fire, and water. He taught that air pos- 
sesses two fundamental properties, heat and dampness. The early 
chemists used the word air in the sense in which the word gas is now 
employed. Thus, we have already learned that hydrogen was first 
called inflammable air. 

The terms atmosphere and air are often used inter- 
changeably, though by air we usually mean a limited por- 
tion of the atmosphere. Many skillful chemists have 
studied the action of air on living things, its relation to 
combustion, the effect of its weight, its composition, and 
its varied properties. Their work has contributed many 
fundamental facts to science. 

General Properties of the Atmosphere. Air has 

weight. We often use the expression " light as air." But 
a cubic foot of air weighs 1.28 oz. and a room 40 x 50 X 25 
ft. contains about two tons of air. The total weight of the 
atmosphere has been estimated to be five thousand millions 
of millions of tons. This enormous mass resting upon the 
earth exerts a pressure which is about fifteen pounds on 
every square inch. This amount of pressure upon a 


62 Descriptive Chemistry. 

square inch is called "an atmosphere," and it is some- 
times used as a unit of pressure. Thus, three atmospheres 
means a pressure of forty-five pounds per square inch. It 
is this pressure which causes water to rise in pumps and 
flow through siphons. Atmospheric pressure is exerted 
in all directions and is variable. It is measured by the 
barometer. The normal or standard pressure of the at- 
mosphere is equal to the weight of a column of mercury 
one square inch in cross section and 29.92 in. high, or one 
square centimeter in cross section and 760 mm. high. But 
since atmospheric pressure is at the rate of fifteen pounds to 
the square inch, it is necessary to know the height only 
of the mercury column in order to know the pressure. 

The pressure of the atmosphere varies as the height and the compo- 
sition of the atmosphere vary, and the barometer changes accordingly. 
The weight of a liter of dry air at o and 760 mm. is i .293 gm. 

The appreciable movements of the atmosphere are the winds. 

Ingredients of the Atmosphere. The atmosphere is a 
mixture of several gases. But since this mixture always 
contains about 78 parts of nitrogen and 21 parts of oxygen 
by volume, we often speak of air as consisting solely of 
these two gases. Besides this large proportion of oxygen 
and nitrogen, the air always contains small and variable 
proportions of water vapor and carbon dioxide gas. Be- 
sides these four ingredients, air always contains the gases 
argon and helium, and usually ozone, hydrogen, hydrogen 
peroxide, compounds related to ammonia and nitric acid, 
dust, and germs. The composition varies but slightly in 
different localities. Near the city air may contain a rela- 
tively larger proportion of dust, ammonia, sulphur com- 
pounds, and acids ; in the country the proportion of ozone 
is relatively large ; at the ocean the air contains consider- 
able salt. 

The Atmosphere Nitrogen. 63 

General Properties of Nitrogen. The chemical ele- 
ment, nitrogen, constitutes about 78 per cent of the atmos- 
phere (by volume). It is a colorless gas, and has no taste 
or odor. It is somewhat lighter than air, and is very 
slightly soluble in water. In many respects it differs 
markedly from oxygen. Thus it will not support combus- 
tion, neither will it burn nor sustain life. Animals die if 

left in nitrogen. 


The fact that a candle flame quickly goes out and a mouse soon dies 
in nitrogen was first observed by Rutherford, an English physician, 
who discovered the gas in 1772. Soon after, Lavoisier showed the true 
relation of nitrogen to the atmosphere. To emphasize the inability 
of the gas to support life, he called the new gas azote, the name now 
used for it by some French chemists. 

Nitrogen is not poisonous, for a large proportion of the 
air we breathe is nitrogen. Its function in the atmosphere 
is to dilute the oxygen. It is an inert element. It com- 
bines with only a few other elements, and many of its 
compounds easily decompose. 

Oxygen and Nitrogen in the Atmosphere. The chem- 
ical activity of the atmosphere is due to the free oxygen 
it contains. We have already learned that oxygen is an 
i active chemical element. If the air were largely oxygen, 
rusting and decay would proceed with astounding rapidity, 
and fires once started would burn with .great violence. On 
kthe other hand, nitrogen is inactive. And if the air con- 
tained much more than the normal amount, chemical 
action would be slower. Oxygen alone is too active, 
while nitrogen alone is inactive. To be serviceable to 
man, oxygen must be diluted with nitrogen, while nitro- 
gen must be accompanied by a small proportion of 

64 Descriptive Chemistry. 

The presence of oxygen and nitrogen in the atmosphere, and the 
functions of the two gases, were first clearly explained by Lavoisier in 
1 777, though many others Boyle, Priestley, Rutherford, and Scheele 
helped solve the problem. 

Composition of the Atmosphere. Samples of air from 
various parts of the globe show a remarkable uniformity 
of composition. Until 1895 it was supposed that pure air 
consisted solely of oxygen and nitrogen. But it has been 
found that about one per cent of the gas hitherto called 
nitrogen is argon, a gas so much like nitrogen, and so 
difficult to separate from the latter, that for years it had 
been overlooked (see Argon, below). According to the 
most recent results, the following is 




By volume. 

By weight. 






Argon .... 


The composition of the atmosphere was studied by Priestley, but his 
results were conflicting. Cavendish, in 1781, was the first to show that 
the proportion of oxygen and nitrogen in air is nearly constant. Since 
his time this result has been confirmed by many chemists, especially by 
Bunsen, who is widely known as the inventor of the Bunsen burner, 
which is used as a source of heat in chemical laboratories. 

The Volumetric Composition of the Air may be found 
by introducing a known volume of pure air into a eudiom- 
eter and exploding it with a known volume of hydrogen. 
The oxygen of the air combines with twice its volume of 
hydrogen, forming a minute quantity of water ; hence one 

The Atmosphere Nitrogen. 

third of the diminution in volume is the volume of oxygen 
in the air. The difference between the volume of oxygen 
found and the original volume of air is the volume of 

An illustration will make this experiment clear. Suppose (i) we 
mix and explode loocc. of air and 50 cc. of hydrogen, or 15000. in all, 
and (2) that the residue measures 87 cc. Now, 150 87 = 63, hence 
63 cc. of the total volume combined to form water. But one third of 
63 cc. is oxygen, which came from the original volume of air. Hence, 
63 -r- 3 = 21, the volume of oxygen in 100 cc. of air. The remainder, 
79 cc., is nitrogen, argon, 
and other gases. 

Another Method, <;often 
used to determine the volu- 
metric composition of the 
air, is based on the fact 
that phosphorus will com- 
bine slowly with oxygen, 
even at the ordinary tem- 
perature. The operation is 
performed in an apparatus 
like that shown in Figure 1 3. 
A piece of phosphorus, C, 
attached to a wire, is 
inserted into a graduated 
glass tube, />, containing a 

measured volume of air. 
White fumes indicate im- 
mediate action. These 
fumes are solid particles 
of an oxide of phosphorus 

FIG. 13. Apparatus for determining the com- 
position of air by phosphorus. 

called phosphorus pentoxide. 'They soon dissolve in the water, which 
rises higher in the tube, as the oxygen combines with the phosphorus. 
In a few hours the phosphorus is removed, and the volume of gas is 
read. The difference between the first and last volumes is oxygen. The 
gas remaining in the tube is, of course, a mixture of nitrogen and argon. 
In performing this experiment unusual care must be taken not to touch 
the phosphorus with the bare hands. 

66 Descriptive Chemistry. 

The Gravimetric Composition of Air was first accurately 
determined in 1841 by the French- chemists, Dumas and 
Boussingault. The average result of many experiments 

tTT'O o Lm 

Oxygen . . . . 23 parts by weight. 
Nitrogen ... 77 parts by weight. 

We know, however, that the correct proportions are 

Oxygen .... 23.2 parts by weight. 
Nitrogen . . . 75.5 parts by weight. 
Argon .... 1.3 parts by weight. 

They passed pure air through a weighed tube containing copper, and 
arranged so that heat could be applied. The oxygen of the air com- 
bined with the copper, while the nitrogen passed on into a weighed globe. 
Both tube and globe increased in weight. The increase in the tube was 
the weight of the oxygen, while the increase in the globe was the weight 
of the nitrogen. 

Water Vapor in the Atmosphere. Water vapor is 
always present in the atmosphere, owing to the constant 
evaporation from the ocean and other bodies of water. 
The total amount present is large, though variable. A 
given volume of air will absorb a definite volume of water 
vapor and no more, and the amount depends largely upon 
the temperature. Air containing its maximum amount of 
water vapor is said to be saturated at that temperature, or 
to contain 100 per cent of water vapor. The saturation 
point is also called the dew point. On a pleasant day the 
relative humidity of the air, i.e. the amount of water 
vapor present, may vary from 30 to 90 per cent, the aver- 
age being about 50 per cent. Warm air holds more vapor 
than cool air. The amount of water vapor in the air has 
a marked influence on the physical condition of man. 
The depressing weather during " dog days " is due to the 

The Atmosphere Nitrogen. 67 

high relative humidity of the air, which sometimes reaches 
95 per cent. The absence of life in deserts is largely due * 
to the dry air .above them. Much of the languor felt in a 
" close " room or crowded hall is partly caused by the 
excess of water vapor in the "bad" air. The presence of 
water vapor in the air is shown by the moisture which col- 
lects on the outside of a vessel containing cold water, such 
as a pitcher of iced water. The moisture comes from the 
air around the vessel. For a similar reason, water pipes 
in a cellar and the cellar walls themselves are moist in 
summer. The deliquescence of calcium chloride, common 
salt, and other substances likewise reveals the presence of 
water vapor in the air (see Deliquescence). 

When the temperature of the air falls, the water vapor condenses and 
is deposited in the form of dew, rain, fog, mist, frost, snow, sleet, or 
hail. The clouds are masses of water vapor which has been condensed 
by the cold upper air. 

Carbon Dioxide in the Atmosphere. Carbon dioxide 
is one product of the respiration of animals, and of the 
combustion and decay of organic substances. By these 
processes an immense quantity of carbon dioxide is being 
constantly poured into the atmosphere. The quantity in 
the atmosphere is variable, though not between such wide 
limits as the water vapor. The proportion in normal air 
is about 4 parts in 10,000 parts of air. Over the ocean 
the proportion is smaller, but in the air of cities it is 
greater. In crowded rooms the proportion is often as 
high as 33 parts in 10,000, because carbon dioxide is 
exhaled faster than it can be removed. The proportion of 
carbon dioxide in the atmosphere as a whole is practically 
constant, largely owing to the fact that this gas is an 
essential food of plants (see Carbon Dioxide). The pres- 
ence of carbon dioxide in the air is detected by limewater. 

68 Descriptive Chemistry. 

If Hmewater is exposed to the air, the carbon dioxide unites with the 
lime in the limewater, forming a thin, white crust of insoluble calcium 
carbonate on the surface of the Hmewater. If air is drawn through lime- 
water, the liquid becomes milky, because the particles of calcium carbon- 
ate are suspended in the liquid. The purity of air is often determined 
by finding out what proportion of carbon dff>xide it contains. If a 
known volume of dry air is drawn through a known weight of Hmewater 
or similar liquid, the increase in weight will be the weight of carbon 
dioxide in the volume of air used. 

The different gases in the atmosphere are not arranged 
in layers according to their densities. They are in con- 
stant circulation (see Diffusion). Hence carbon dioxide, 
though heavier than oxygen and nitrogen (volume for vol- 
ume), does not remain nearest the ground, but is distrib- 
uted through the air. In a few exceptional localities, 
carbon dioxide arises from volcanoes faster than it can 
diffuse, and fills the adjacent valley. 

Argon in the Atmosphere. Argon is a colorless, odor- 
less gas. Its chief characteristic is its chemical inactivity. 
No compounds of argon have as yet been prepared or 
discovered. The name argon is happily chosen, being 
derived from Greek words signifying inert. It constitutes 
0.94 per cent by volume of the atmosphere, or 1.3 per 
cent by weight. 

Argon was discovered in 1894 by Rayleigh and Ramsay. Rayleigh 
had found that nitrogen from air weighed more than an equal volume 
of nitrogen obtained from compounds of nitrogen. Consequently, they 
believed that the nitrogen from air contained another gas hitherto over- 
looked. A series of elaborate experiments showed that after all the 
oxygen and nitrogen was removed from purified air, there still remained 
a small quantity of a new gas, which they called argon. It may be pre- 
pared (i) by passing pure air over healed copper to remove the oxygen, 
and then the remaining gas over heated magnesium or calcium to remove 
the nitrogen ; or (2) by passing electric sparks through a mixture of air 
and oxygen, and removing the compound of oxygen and nitrogen as fast 

The Atmosphere Nitrogen. 69 

as it is formed. The latter method is a repetition of the one used by 
Cavendish when he determined the composition of air, and he would 
have no doubt discovered argon had he continued his investigations. 

Inert Gases in the Atmosphere. Helium, neon, krypton, and xenon 
have recently been discovered by Ramsay. At present little is known 
about these gases. They resemble argon in being inactive chemical 
elements. They constitute an exceedingly minute proportion of the 
atmosphere. Helium is also found in certain rare minerals, in the gases 
from some mineral springs, and in the atmosphere of the sun. It is 
about twice as heavy as hydrogen. According to Ramsay, " it is prob- 
able that helium is continually escaping from the earth in small quantities 
in certain regions.' 1 

Air is a Mixture, in spite of the fact that we speak of 
its "composition." Chemical compounds have two invari- 
able characteristics : viz., (i) their components are in a fixed 
proportion, and (2) their formation and decomposition are 
usually attended by definite evidences of chemical action, 
such as light, heat, change of color and form, etc. The 
following facts show that air is a mixture of free gases : 

(1) The proportion of oxygen and of nitrogen is not 
fixed, but varies between small limits, which may be 
detected by accurate analysis. 

(2) When nitrogen and oxygen are mixed in the propor- 
tions which form air, the product is exactly like air, but 
the act of mixing gives no evidence of chemical action. 

(3) When air is dissolved in water, a greater proportion 
of oxygen than of nitrogen dissolves. If the oxygen and 
nitrogen were combined in the air, the dissolved air would, 
of course, have the same composition as air itself. 

Liquid Air is a mixture of the liquefied gases which con- 
stituted the air used. It is a milky liquid, owing to the 
presence of solid carbon dioxide and ice. If these solids 
are removed by filtering, the filtrate has a pale blue tint. 
It is slightly heavier than water. It is intensely cold, its 

jo Descriptive Chemistry. 

temperature being about 200 C. It boils at about 
190 C. under atmospheric pressure. If a tumbler is 
filled with liquid air, the latter boils vigorously, the sur- 
rounding air becomes intensely cold, frost gathers on the 
tumbler, and in a short time the liquid air will have 

entirely disappeared into 
the air of the room. If, 
however, the liquid air is 
placed in a Dewar's bulb 
or flask, it evaporates so 
slowly that some will remain 
in the flask several hours. 

The Dewar's bulb (Fig. 14) 
consists of two flasks, one within 
the other, attached at the top ; the 
space between the flasks is a 
vacuum. Sometimes the outer 
surface of the inner flask is coated 
with mercury or silver, which 
helps to protect the liquid air from 
the heat of the atmosphere. In 
transporting liquid air a large 
Dewar's bulb or similar device is 
FIG. 14. A Dewar's bulb. used. One form consists of a 

large metal can wrapped with 

many thicknesses of felt and slipped into a larger can covered with 
canvas or felt. The liquid air is put in the inner can and a loose 
stopper or piece of felt is placed over the mouth. The liquid may also 
be kept in these cans for some time with only a moderate loss, unless 
the surrounding temperature is exceptionally high. 

Liquid air, owing to its extremely low temperature, pro- 
duces remarkable physical changes. A tin or iron vessel 
which has been cooled by liquid air is so brittle that it may 
often be crushed with the fingers. Nearly all plastic or 
soft substances, including many kinds of food, when im- 

The Atmosphere Nitrogen. 71 

mersed in liquid air, become hard and brittle, leather being 
the only important exception. Mercury freezes so hard in 
liquid air, that it may be used as a hammer to drive a nail. 
When liquid air is put in a teakettle standing on a block 
of ice, the liquid air boils vigorously. If the kettle of 
liquid air is placed over a lighted Bunsen burner, frost and 
ice collect on the bottom of the kettle, because the intense 
cold of the kettle solidifies the water vapor and carbon 
dioxide, which are the two main products of burning 
illuminating gas. If water is now poured into the kettle, 
the liquid air boils over and the water is instantly frozen ; 
the water is so much hotter than the liquid air that the latter 
boils more violently, and since its rapid evaporation causes 
absorption of heat, the water gives up its heat and becomes 
ice. Ordinary liquid air* is from one half to one fifth liquid 
oxygen, and will support combustion. A red-hot rod of 
steel or of carbon burns brilliantly in this cold liquid. 

Numerous applications of liquid air have been proposed, but thus far 
they have not passed the experimental stage. It has been suggested 
that it be used as a refrigerant instead of ice, for ventilating and cooling 
rooms, as a blasting material, for removing diseased flesh from a wound, 
for destroying refuse, and as a commercial source of oxygen. The last 
use is based primarily on the fact that as liquid air evaporates, the 
nitrogen passes off first, and in a short time relatively pure oxygen 
remains (see Oxygen). 

A little liquid air was produced in 1883 with considerable labor and 
at an enormous expense. Now it is Easily manufactured in large quan- 
tities at a comparatively low cost. In the older methods of preparing 
liquefied gases, the gas was subjected to tremendous pressure and a low 
temperature. At present, air is liquefied by a different method. Com- 
pressed air cooled by water is forced through a pipe with a small open- 
ing into a larger cylinder called the liquefier. As it escapes into the 
liquefier it expands and its temperature falls, because expansion is a 
cooling process. The temperature of the liquefier is thus reduced, so 
that the air, which continues to enter, expands at such a low temperature 
that it becomes a liquid. 

72 Descriptive Chemistry. 


Occurrence. Nitrogen, besides comprising four fifths 
of the atmosphere, is a component of nitric acid and am- 
monia, and of the many compounds related to them. It 
is also an essential constituent of animal and vegetable 

The name nitrogen was given to the gas by Chaptal from the fact 
that it is a component of niter, an old name of potassium nitrate. 

Preparation. Nitrogen is usually obtained from the air by remov- 
ing the oxygen by phosphorus. A tall jar is placed over burning 
phosphorus contained in a shallow dish floating in a large vessel of 
water. The oxygen combines with the phosphorus, leaving nitrogen, 
more or less pure, in the jar. Other methods may be used, such as 
decomposing ammonium nitrite by heat, or passing air over heated 

Additional Properties. In addition to its inertness, already men- 
tioned, nitrogen is a little lighter than air, and is very sparingly soluble 
in water. Its density is 0.972 (air = i). One liter at o C. an'd 760 mm. 
weighs i.256gm. One hundred liters of water dissolve only 1.5 1. at the 
ordinary temperature. It combines with magnesium and a few other 
metals at a red heat, forming nitrides. Electric sparks cause nitrogen to 
combine with oxygen and with hydrogen, forming ultimately nitric acid 
and ammonia, hence these substances or others related to them are 
often found in the rain which falls during a thunder storm. 

Relation of Nitrogen to Life. Oxygen, carbon diox- 
ide, and water vapor are essentially related to the life of 
plants and animals. Nitrogen is also vitally connected 
with different forms of life. Atmospheric nitrogen merely 
dilutes the oxygen. Although we live in an atmosphere 
containing such a large proportion of nitrogen, we cannot 
assimilate it. According to a reliable authority, " the air 
as it leaves the lungs contains 79.5 per cent of nitrogen," 
and hence cannot become a part of the body. Yet all flesh 
contains nitrogen, and the rejected waste products of ani- 

The Atmosphere Nitrogen. 73 

mals are largely combined nitrogen. The nitrogen needed 
by animals must be in combination to become available. 
And it is taken in the form of nitrogenous food, such as 
lean meat, fish, wheat and other grains. 

Most plants take up combined nitrogen from the soil in 
the form of nitrates (compounds derived from nitric acid) 
or of ammonia. Hence combined nitrogen is being con- 
stantly taken from the soil, and in order to preserve the 
fertility of the soil, nitrogen must be supplied. This is done 
by allowing nitrogenous organic matter to decay upon the 
soil, or by adding to the soil a fertilizer, which is a 
mixture containing nitrogen compounds. Recently it 
has been shown that leguminous plants, such as peas, 
beans, and clover, take up nitrogen from the air by means 
of bacteria, which are in nodules on their roots. 


1. What is the atmosphere? What is air? What is the literal 
meaning of the word atmosphere? What is the wind? 

2. Develop the topics: (a) atmospheric pressure, (b) occurrence of 
nitrogen, (c) volumetric composition of the air, (W) gravimetric com- 
position of the air, (e) water vapor in the atmosphere, (/") carbon 
dioxide in the atmosphere, (g) air is a mixture. 

3. Define and illustrate the terms : (#) an atmosphere, (<) normal 
pressure, (c) standard pressure, (d) dew point, (e) relative humidity, 
(/) inert. 

4. What are the two chief ingredients of the atmosphere? The per- 
manent ingredients ? The variable ingredients ? The ingredients found 
in traces? What are sometimes found in the air of cities? 

5. What is the symbol of nitrogen? What are its general proper- 
ties? Its special properties? What is its main function in the atmos- 
phere? How may it be prepared? 

6. When and by whom was nitrogen discovered? Why and by 
whom was it named "azote 11 and "nitrogen 11 ? 

7. What is the relation of nitrogen to animal and to vegetable life? 

74 Descriptive Chemistry. 

8. Compare the functions of oxygen and nitrogen in the atmos- 
phere. What famous chemists helped solve this problem? 

9. State the composition of pure air (a) by volume, and (b) by 

10. Give a brief biographical account of (a) Cavendish, (^) Dumas, 
(c) Rutherford. (See Appendix, 4.) 

11. What is a cloud? The dew? Why does moisture gather on 
cellar walls? Why are mines often damp? What is (a) rain, (t>) fog, 

(c) mist? 

12. Describe the action of air upon (a) limewater, and (b} calcium 

13. How does the atmosphere illustrate the diffusion of gases? 

14. What is argon? Give a brief account of (a) its discovery, () its 
properties, (c) its method of preparation. What proportion of pure 
air is argon? What is the significance of the name argon f 

15. Give a brief account of helium, neon, krypton, and xenon. 

1 6. What is liquid air? What are its chief properties? State 
briefly its method of manufacture. Describe its action (a) upon solids, 
such as rubber, (b) upon liquids, such as mercury, (c) upon hot steel, 

(d) when evaporated quickly. Describe a Dewar's bulb. 


1. If a man inhales 18 cu. ft. of air an hour, what weight of oxy- 
gen does he consume in 24 hr. ? 

2. What is the weight of air in a room, 6x6x3111., if a liter of 
the air weighs 1.3 gm. ? 

3. A mixture of 25 cc. of air and 50 cc. of hydrogen is exploded. 
The residue measures 60.3 cc. What per cent of oxygen did this 
sample of air contain ? 

4. How many kilograms of pure air are needed to yield 100 kg. of 
oxygen ? 

5. Express in inches the following barometer readings : (a) 760 
mm., (<) 740 mm., (c) 75 cm., (d) 0.749 m., (e) 7.67 dm. 

6. Dumas and Boussingault, in 1841, found in a sample of air, 
12 -373 g m - of nitrogen and 3.68 gm. of oxygen. What per cent of 
each was found? 

7. What is the weight at o C. and 760 mm. of (#) 1000 cc. of dry 
air? Of () 750 1., (c) 1750 cc., (d) 850 cu. m.? 



Law and Theory. We discover facts by observation 
and experiment. Facts which always oc"cur under the 
same circumstances soon become well established. Such 
facts are often -summarized in a brief statement called a 

Sometimes the word law is used in the sense of the uniform behavior 
summarized in the brief statement. Hence, in a narrow sense, a law 
is a statement of a fact, but in a broad sense a law is the fact itself. 
Thus, the law of definite proportions (soon to be discussed) is either 
(i) a brief statement of the general fact of definite proportions of ele- 
ments in compounds, or (2) the uniform behavior itself as far as the 
composition of chemical compounds is concerned. 

The cause of many scientific facts is unknown. The 
explanation we give, or the statement we make, of the 
cause of facts is called a theory. Laws are statements of 
fact, theories are statements of the supposed cause of facts. 
Thus we know that chemical compounds have a definite 
composition, because we have discovered by experiment 
the facts on which this law is based ; and we have framed 
a theory, which, as far as our present knowledge is. con- 
cerned, is a satisfactory explanation of the cause of the 
general fact of definite composition. Laws seldom change, 
but theories are often modified. Laws are the result of 
experiment, theories are the outcome of mental operations. 


76 Descriptive Chemistry. 

We accept a certain theory until a more satisfactory one is 
proposed. If a fact is not well established or is not gen- 
eral, we account for it by an hypothesis. An hypothesis 
is a guess or supposition concerning the cause of some 
particular fact or set of facts, and it is usually proposed as 
a basis for making further experiments. Hypotheses often 
lead to theories. 

Laws, theories, and hypotheses are of great service in 
chemistry, since they help us gather into intelligible state- 
ments a vast number of facts which are apparently not 
related. They also assist in discovering facts. 

Law of Definite Proportions by Weight. When the 
metal magnesium is heated in the air, it burns with a 
dazzling flame into a grayish powder, due to combination 
with oxygen. If a known weight of magnesium is heated 
in a crucible, so that the product cannot escape, a remark- 
able relation is revealed. In order to burn completely 1.5 
gm. of magnesium, i gm. of oxygen is necessary; and 
the product, magnesium oxide, weighs 2.5 gm. This 
product contains, therefore, 60 per cent magnesium and 
40 per cent oxygen. Accurate repetitions of this experi- 
ment have shown that this proportion by weight is fixed 
and definite. Again, if all the oxygen is driven from a 
weighed quantity of potassium chlorate by heating this 
compound in a crucible, 39.18 per cent of oxygen is 
always obtained. This means that the proportion of 
potassium, chlorine, and oxygen which makes up potas- 
sium chlorate is fixed and definite. Otherwise, the prop- 
erties of potassium chlorate would vary. Experiments 
similar to these show that in all chemical compounds the 
different components are always present in a definite and 
unvarying proportion by weight. There are no exceptions 
to this general fact. This constancy of proportion in 

Law of Multiple Proportions. 77 

chemical compounds is stated as the Law of Definite Pro- 
portions by Weight, thus : - 

A given chemical compound always contains the same 
elements in the same proportions by weight. 

Sometimes it is condensed into this form : 

A chemical compound has a definite composition by weight. 

This law is one of the fundamental laws of chemistry. It is so firmly 
believed that if the composition of a compound is found by analysis to 
vary, chemists conclude that the experimental work is incorrect or that 
the compound is impure. The law was established as the outcome of 
a controversy between two French chemists, Proust (1755-1826) and 
Berthollet (1748-1822). The discussion lasted from 1799 to 1806. 
Berthollet believed that compounds might have a varying composition. 
Indeed, by his experiments he detected " gradual changes " in com- 
position. But Proust showed that Berthollet analyzed mixtures and 
not compounds. In a mixture the parts may be present in any propor- 
tion. Subsequent experiments have only strengthened our confidence 
in this law. 

Law of Multiple Proportions. Proust showed that 
some elements combine in more than one proportion, 
and thereby produce distinct compounds. But he failed to 
notice that if the weight of one element is constant, the 
varying weights of the other element are in a simple mul- 
tiple relation to each other. Dalton discovered this gen- 
eral fact about 1804. The composition of compounds is 
usually expressed in per cent ; but such expressions in a 
series of compounds reveal nothing about multiple rela- 
tions. If, however, a constant weight is adopted as a unit 
for one component, and the composition of the series of 
compounds is expressed in terms of this unit, then the 
simple multiple relation which exists between the weights 
of the other component is clearly seen. Thus, we learn> 
little from the statement that the two compounds of carbon 

Descriptive Chemistry. 

and oxygen contain 73 and 57 per cent of oxygen. But 
if in expressing the composition of these compounds 
we adopt 12 as the weight of carbon, the weights of 
oxygen become 32 and 16, i.e. the weights of oxygen are 
simple multiples. The five compounds of oxygen and 
nitrogen, which will soon be studied, aptly illustrate this 
fact : 








Nitrogen. Oxygen. 



;en. Oxygen. 

Nitrous oxide .... 

63.6 36,4 




Nitric oxide 

j.6 6 zi A. 




Nitrogen trioxide . . . 

36.8 63.2 




Nitrogen peroxide. . . 

30.4 69.6 




Nitrogen pentoxide . 

25.9 74.1 




From this table it is clear that the weights of oxygen 
combined with the same weight of nitrogen are as 1:2: 
3:4:5, i.e. they are simple multiples of each other. 

The general fact of multiple proportions is expressed in 
the Law of Multiple Proportions, thus : - 

When two or more elements unite to form a series of 
compounds, a fixed weigJit of one element so combines with 
different weights of the other element that the relations be- 
tween the different weights can be expressed by small whole 

This law, like the law of definite proportions, is a fun- 
damental law of chemistry, and together they have pro- 
foundly influenced its theoretical and practical progress. 




The Atomic Theory. 79 

The Atomic Theory of the constitution of matter was 
proposed by Dalton to explain the laws of definite and 
multiple proportions. This theory assumes (i) that the 
chemical elements consist ultimately of a vast number of 
very small, indivisible particles or atoms, (2) that the 
atoms of the same element have the same weight, (3) that 
atoms of different elements have different weights, and (4) 
that chemical action is union or separation of the atoms of 
the elements. 

Let us now consider how this theory explains the facts 
summarized in the laws of definite and multiple propor- 
tions, (i) When magnesium combines with oxygen, 1.5 
parts by weight of magnesium combine with one part by 
weight of oxygen. Analysis of the product magnesium 
oxide shows that this proportion is constant; that is, 
pure magnesium oxide always contains the elements mag- 
nesium and oxygen in this proportion. Now, according to 
the atomic theory, magnesium oxide is the product of the 
union of indivisible atoms of magnesium and indivisible 
atoms of oxygen. It therefore follows that when magne- 
sium and oxygen unite, atom for atom, the magnesium 
oxide must contain the two elements in the proportion of 
the weights of their atoms, i.e. it must always have the 
same composition. It is immaterial whether the actual 
weights of these elements which combine are in the pro- 
portion of i to 1.5, because whatever is in excess of this 
proportion will be left uncombined. For example, if we 
start with i gm. of oxygen and 2 gm. of magnesium, then 
0.5 gm. of magnesium will be left uncombined. Thus the 
atomic theory explains the law of definite proportions. (2) 
But atoms do not always combine in the simple proportion 
of i to i. They may combine in the proportions of i to 2, 
2 to 3, i to 3, i to 4, etc. But according to the atomic 

8o Descriptive Chemistry. 

theory atoms are assumed to be indivisible. Hence, if we 
assume the atomic theory, the proportions of the weights 
of different elements in a series of compounds must be 
simple proportions, i.e. the elements must unite in accord- 
ance with the law of multiple proportions. To illustrate : 
There are two compounds of carbon and oxygen. Since 
atoms are indivisible, the simplest combinations of the atoms 
are (i) one atom of carbon to one atom of oxygen, and (2) 
one atom of carbon to two atoms of oxygen. Analysis 
shows that in the first compound the proportion of carbon 
to oxygen is 6 to 8. According to the theory, the propor- 
tion in the second compound should be 6 to 16; this pro- 
portion is verified by analysis. In other words, if we 
adopt 6 as the weight of carbon in its two oxides, then the 
weights of oxygen are in the simple proportion i to 2. 

Atoms and Molecules. It should not be forgotten that 
the laws of definite and multiple proportions deal with 
facts, and that the atomic theory deals with conceptions 
which may be true, but which cannot be proved to be 
true. We often speak of atoms as if they could be per- 
ceived by the senses, but we do so simply because such 
expressions help us describe, study, and interpret chemical 
action. According to the present views, atoms do not, as 
a rule, exist in the uncombined state. As soon as atoms 
are freed from combination, they at once unite with some 
other atom or atoms. The smallest particle of matter 
which can exist independently is not, therefore, an atom, 
but a group or combination of atoms. These groups of 
atoms are called molecules. If the atoms in a molecule 
are atoms of the same element, then the molecule is a 
molecule of an element; but if the atoms of different 
elements are combined, then the molecule is the molecule 
of a compound. All matter, as a rule, consists of mole- 

Chemical Symbols. 8 

cules, and the molecules are made up of atoms. A mole- 
cule of a few elements contains only one atom. Chemists 
define a molecule as the smallest part of a compound or 
of an element which can exist in the free state and mani- 
fest the properties of the compound. Thus, the smallest 
particle of water is a molecule of water, but a molecule of 
water contains smaller particles still, viz., atoms of hydro- 
gen and oxygen. We may define an atom as the indivis- 
ible constituent of a molecule. It is also the smallest 
particle of an element which takes part in chemical 

Our views regarding molecules are based on extensive study of the 
physical properties of gases. The molecule is often spoken of as the 
physical unit, because in physical changes molecules are not decomposed. 
Whereas the atom is the chemical unit, because it enters into all chemi- 
cal action. The molecule is chemically divisible, but the atom is 
chemically indivisible. 

Chemical Symbols, which were mentioned in Chapter I, 
are designed to represent single atoms. Thus, H repre- 
sents one atom of hydrogen, O one atom of oxygen, N one 
atom of nitrogen. If more than one atom is to be desig- 
nated, the proper numeral is placed before the symbol, 

2 H means 2 atoms of hydrogen. 

3 O means 3 atoms of oxygen. 

4 P means 4 atoms of phosphorus. 

But if the atoms are in chemical combination, either with 
themselves or with other atoms, then a small numeral is 
placed after and a little below the symbol, thus : 

H 2 means 2 atoms of hydrogen in combination, 
N 3 means 3 atoms of nitrogen in combination, 
P 4 means 4 atoms of phosphorus in combination. 

8 2 Descriptive Chemistry. 

Chemical Formulas. A formula is a group of symbols 
which is designed to express the composition of a com- 
pound. In writing a formula the symbols of the different 
atoms making up the compound are placed side by side. 
Thus, H 2 O is the formula of water, because this group of 
symbols is the simplest expression of the facts which are 
known about this compound. Similarly, KC1O 3 is the 
formula of potassium chlorate. These symbols might be 
written in a different order, but usage has determined the 
order in this, as in most cases. A formula represents one 
molecule. Hence, KC1O 3 represents one molecule of 
potassium chlorate, and means that the molecule of this 
compound contains one atom each of potassium and chlo- 
rine and three atoms of oxygen. If we wish to designate 
several molecules, the proper numeral is placed before the 
formula, thus : 

2 KC1O 3 means 2 molecules of potassium chlorate. 

3 H 2 O means 3 molecules of water. 

4 H 2 SO 4 means 4 molecules of sulphuric acid. 

In certain compounds some of the atoms act like a single 
atom in chemical changes. This fact is often expressed by 
inclosing the group of atoms in a parenthesis, or by sepa- 
rating it from the rest of the formula by a period. Thus, 
the formula of ammonium nitrate is (NH 4 )NO 3 . Simi- 
larly, the formula of alcohol is often written C 2 H 5 . OH, 
because the groups C 2 H 5 and OH act as units. The use 
of the period is confined mainly to organic and mineralogi- 
cal chemistry. It is sometimes omitted, especially if the 
composition of the compound is well understood. If a 
group of atoms is to be multiplied, it is placed within a 
parenthesis. Thus, the formula of lead nitrate is Pb(NO 3 ) 2 . 
This means that the group NO 3 is to be multiplied by 2. 

Chemical Equations. 83 

The formula 2 Pb(NO 3 ) 2 means that the whole formula is 
to be multiplied by 2. 

Symbols and formulas are sometimes used to represent an indefinite 
amount of an element or compound. Thus, O may mean oxygen and 
H.jSO 4 sulphuric acid, regardless of the amount. This use of symbols 
and formulas saves time, but it is not scientific. They are often thus 
used to label bottles in a laboratory. Such a departure from accuracy 
should not be allowed to obscure their real meaning. 

The complete significance ot symbols and formulas can be grasped 
only by their intelligent use. They should not be committed to mem- 
ory slavishly. It is desirable, however, to learn the common ones 
while the substances they represent are being studied, and consider 
their relations more fully when the needed facts have accumulated. 
(See Chapters IX and XIII.) 

A Chemical Reaction is a special or limited chemical 
change. When potassium chlorate is heated, the chemical 
change results finally in the liberation of all the oxygen 
and the formation of potassium chloride. Such a change 
is called the reaction for preparing oxygen from potassium 
chlorate, or the reaction for the decomposition of potas- 
sium chlorate. Obviously, the study of chemistry is 
largely a study of reactions. 

Chemical Equations. In expressing various facts 
about chemical reactions, it is customary to use an equa- 
tion consisting of the proper symbols or formulas. Sub- 
stances entering into the initial stage of a reaction are 
called factors, and those present in the final stage are 
called products. The symbols and formulas of the factors 
connected by the sign plus ( -f- ) are placed at the left of 
the sign of equality, and those of the products at the right. 
Equations are usually read from left to right. Occasion- 
ally the words reaction and equation are used as synonyms, 
but such a use is inaccurate and confusing. 

84 Descriptive Chemistry. 

When magnesium burns in the air or in oxygen, mag- 
nesium oxide is formed. The simplest equation for this 
reaction is 

Mg + O = MgO 
Magnesium ' Oxygen Magnesium Oxide 

This equation is read : Magnesium and oxygen form mag- 
nesium oxide. It means, also, that when magnesium and 
oxygen' react, one atom of magnesium unites with one 
atom of oxygen and forms one molecule of magnesium 
oxide. The simplest equation for the preparation of hy- 
drogen by the reaction of zinc and sulphuric acid is 

Zn+ H 2 SO 4 = H 2 + ZnSO 4 
Zinc Sulphuric Acid Hydrogen Zinc Sulphate 

This equation is read: Zinc and sulphuric acid form (or 
produce) hydrogen and zinc sulphate. It means, further, 
that one atom of zinc and one molecule of sulphuric acid 
form one molecule (or two atoms) of hydrogen and one 
molecule of zinc sulphate. By similar equations we may 
express certain facts about all reactions which are under- 
stood. The above equations might be called ordinary 
chemical equations, or atomic equations. Other forms 
are used, and they will be discussed in Chapters IX, X, 
and XIII. 

The following facts about ordinary chemical equations should be 
noted : 

(1) The sign plus does not necessarily mean addition chemically. 
It does in the equation Mg -f O = MgO, but not in the equation HgO 
= Hg-fO. In the latter the products are merely mixed. The sign 
plus may be expressed by the words and* acted upon, added to, mixed 
with. The sign equality is often read equal, give, form, or produce. 

(2) Equations do not always include all the participating substances. 
In Mg + O = MgO no nitrogen (N) appears because nitrogen takes no 

Exercises. 85 

chemical part in the change, despite the fact that the air is largely 
nitrogen. Similarly, in Zn + H 2 SO 4 = H 2 + ZnSO 4 , no water (H 2 O) 
appears, because the water (in the dilute sulphuric acid) simply serves 
to dissolve the zinc sulphate from the surface of the zinc. A special 
form of equation, called the ionic equation, is used to express chemical 
changes which occur in solution (see Chapter X). 

(3) Equations tell nothing about the heat changes (see Chapter X). 

(4) Most equations represent only the beginning and end of reac- 
tions. Thus, in KC1O 3 = O 3 + KC1 several changes do not appear, 
because the purpose of this equation is to express the complete decom- 
position of potassium chlorate nothing else. 


1. Define law, theory, and hypothesis as used in science. 

2. State the law of definite proportions. Illustrate it. Give a brief 
account of its discovery. 

3. State the law of multiple proportions. Illustrate it. Who dis- 
covered it? When? 

4. State the atomic theory. What are atoms according to this 
theory? How are atoms related to chemical action? How are atoms 
related to molecules? What is a molecule? 

5. What is the symbol of an element? How are they formed? 
Interpret the symbols : H, 2O, N 3 , 2 P, 30, K 2 , S 2 , 2 Cl. 

6. What is the formula of a compound? What does a formula 
represent? Interpret the formulas: H 2 O, 2 H 2 O, KC1O 3 , 4 H 2 SO 4 , 
(NH 4 )NO 3 , C 2 H 5 .OH, Pb(N(X) 2 , Ca(OH) 2 . " 

7. Give the symbols of the following elements : oxygen, hydrogen, 
nitrogen, zinc, copper, magnesium, platinum, iron, sodium, sulphur, 
carbon, mercury. 

8. What elements correspond to the following symbols : Na, Cu, 
K, Zn, S, P, Pt, Pb, H, Hg, Fe, Mg? 

9. Give the formulas of the following compounds : water, potas- 
sium chlorate, sulphuric acid, magnesium oxide. 

10. Define and illustrate the term chemical reaction. 

11. What is a chemical equation ? For what is it used? What are 
factors and products in an equation? How are equations written? 
Illustrate your answer. How are they read ? 

86 Descriptive Chemistry. 

12. Interpret the equation : Mg + O = MgO. 

13. What does the plus ( + ) sign mean in the above equation? 
What other meanings has this sign? 

14. State several facts about equations. 


1. How many centigrams in 1745 kg.? In 250 gm.? In 1425 dg. ? 

2. How many cubic centimeters in 50 1. ? In I cu. dm. ? 

3. What is the weight of (a) loocc. of hydrogen, and (6) 25 1. of 
oxygen, under standard conditions ? 

4. What weight of (a) hydrogen and (<) oxygen can be obtained 
from 1 80 gm. of water ? 

5. What (#) weight and ($) volume of oxygen are necessary to unite 
with 200 kg. of hydrogen ? 

6. What weight of hydrogen is necessary to unite with the oxygen 
in 100 gm. of air to form water ? (Assume that air is one fifth oxygen.) 


Introduction. Many chemical compounds fall naturally 
into one of three groups, long known as acids, bases, and 
salts. Not all compounds, of course, are included in this 
classification. Each group has its characteristic properties, 
'though the groups are closely related and sometimes over- 
lap. Many familiar substances belong to these groups. 
A knowledge of the properties of acids, bases, and salts,, 
of their special behavior, and of their intimate relations is 
essential in the study of chemistry. 

General Properties of Acids, Bases, and Salts. Acids 
have a sour taste. The early chemists detected this 
property, and the word acid (from the Latin acidtts, sour) 
emphasizes the fact. Acids change the color of many 
vegetable substances. Thus, blue litmus is turned red by 
acids. Acids also have the power to decompose most 
carbonates, like limestone, thereby liberating carbon diox- 
ide gas which escapes with effervescence. Most bases 
have a slimy, soapy feeling, and a bitter taste. They turn 
red litmus blue. Caustic soda and ammonium hydroxide 
are bases. Many salts have the well-known salty taste. 
Sodium chloride, the familiar table salt, is an example. 
Usually, they have no action on litmus. 

All acids contain hydrogen, which is usually liberated 
when metals and acids interact. Most acids contain oxy- 
gen. For many years it was thought that oxygen was an 


88 Descriptive Chemistry. 

essential component of all acids, and its name, oxygen 
(derived from Greek words meaning " acid producer ") was 
given by Lavoisier because of this belief (see Discovery of 

We now know that hydrogen, not oxygen, is the 
essential component of all acids. Another necessary 
component of acids is some element like nitrogen, sulphur, 
chlorine, or phosphorus, which belongs to a class of 
elements called non-metals. For this reason it is some- 
times convenient to think of non-metals as the elements 
which form acids. Thus sulphuric acid contains sulphur, 
besides hydrogen and oxygen ; while hydrochloric acid 
contains chlorine, besides hydrogen. 

Bases contain oxygen and usually hydrogen, but their 
distinctive component is a metal, e.g. sodium, potassium, 
calcium. Hence a metal may be properly regarded not 
merely as an element possessing in a varying degree the 
physical properties of hardness, luster, power to conduct 
heat and electricity, but also the chemical property of 
forming bases. 

Salts contain a metal and a non-metal, and most of them 
contain oxygen. Thus, potassium nitrate contains the 
metal potassium and the non-metal nitrogen, besides 
oxygen ; while potassium chloride contains potassium 
and the non-metal chlorine, but no oxygen. 

The nature o*f acids, bases, and salts is clearly shown by 
their chemical relations to each other. When acids and 
bases interact, salts are formed. That is, the acid and 
base destroy more or less completely the marked prop- 
erties of each other and produce a compound which has 
few, and often none, of the properties of the original acid 
or base. The acid and base neutralize each other. An 
example will make this point clear. When hydrochloric 

Acids, Bases, and Salts. 89 

acid and sodium hydroxide interact, sodium chloride and 
water are formed. The chemical change may be written 


HC1 + NaOH. NaCl + H 2 O 

Hydrochloric Acid Sodium Hydroxide Sodium Chloride Water 

This equation represents the facts which have been 
repeatedly verified by experiment. This series of chemi- 
cal changes is called neutralization, and later it will be 
more fully discussed. Taking this equation as a type of 
the chemical changes which occur in neutralization, it is 
clear that in such changes, generally speaking (i) the metal 
of the base takes the place of the hydrogen of the acid, 
thereby forming a salt, while (2) the hydrogen of the acid 
combines with the hydrogen and oxygen of the base to 
form water. In neutralization the hydrogen and oxygen 
of the base act as a unit. This group of atoms (OH) is 
called hydroxyl. Compounds containing this group are 
called hydroxides. Hydroxyl does not exist free and 
uncombined like elements and compounds, but it acts like 
a single atom in many changes. It is called a radical. 
To emphasize the fact that it is a unit, the hydroxyl group 
is sometimes put in a parenthesis, e.g. Ca(OH) 2 . 

Hydroxides are often said to be founded on the water type. Thus 
we have 

Water HOH 

Sodium hydroxide . . . . ' NaOH 
Potassium hydroxide .... KOH 
Calcium hydroxide .... Ca(OH) 2 

Hence we may regard sodium hydroxide and potassium hydroxide 
as water in which the hydrogen atom has been replaced by a metallic 

The words hydroxide, hydrate, and hydroxyl are all derived from 
hudor, the Greek word for water. 

90 Descriptive Chemistry. 

The most characteristic property of acids and bases is, 
then, this power to neutralize each other and thereby form 
salts and water. 

Acids. The common acids are sulphuric acid, hydro- 
chloric acid, nitric acid, and acetic acid. Many acids are 
liquid, as sulphuric and nitric ; a few are gases, as hydro- 
chloric ; others are solid, as tartaric, citric, oxalic. Most 
are soluble in water, and such solutions are familiarly 
called acids. These solutions may be dilute or concen- 
trated, and the general properties vary somewhat with the 
strength. Concentrated acids are usually corrosive and 
should be handled with precaution, even when one is 
thoroughly familiar with their properties. Substances 
which turn blue litmus to red are said to contain an acid, 
to be acid, or to have an acid reaction. The exact nature, 
however, of such a substance must be determined by 
additional tests. 

Many familiar substances are acids or contain them. 
Vinegar, pickles, and similar relishes contain dilute acetic 
acid. Lemon juice is mainly citric acid. Sour milk con- 
tains lactic acid. Unripe fruits, sour bread, and sour 
wines contain acids. " Soda water " is a solution of 
carbonic acid (or more accurately carbon dioxide), and 
" acid phosphate" is a solution of a sour calcium phosphate. 

No brief, satisfactory definition of an acid can be given, 
for chemists do not agree on this point. We might say, 
however, that an acid is a compound containing hydrogen 
which can be replaced by a metal; but this definition 
includes water, since its hydrogen is readily replaced by 
sodium. Not only must the hydrogen of an acid be 
replaced by a metal, but one product of the reaction must 
be a salt. The replacing metal may, of course, come from 
a compound, e.g. an oxide, hydroxide, or carbonate. 

Acids, Bases, and Salts. 91 

Nomenclature of Acids. Oxygen is a component of 
most acids, and the names of these acids correspond to 
the proportion of oxygen which they contain. The best 
known acid of an element usually has the suffix -ic, e.g. 
sulphuric, nitric, phosphoric. If an element forms another 
acid, containing less oxygen, this acid has the suffix -ous, 
e.g. sulphurous, chlorous, phosphorous. Some elements 
form an acid containing less oxygen than the -ous acid ; 
these acids retain the suffix -ous, and have, also, the prefix 
hypo-, e.g. hyposulphurous, hypophosphorous, hypochlo- 
rodl. Hypo- means under or lesser. If an element forms 
an acid containing more oxygen than the -ic acid, such an 
acid retains the suffix -ic, and has, also, the prefix per-, e.g. 
persulphuric, perchloric. The prefix per- means beyond 
or over. The few acids which contain no oxygen have 
the prefix hydro- and the suffix -ic, e.g. hydrochloric, 
hydrobron\i, hydrofluoric. It should be noticed that 
these suffixe^ are not always added to the name of the 
element, but often to some modification of it. 

The nomenclature of acids is well illustrated by the series of chlorine 
acids : 





HC1O 2 


HC10 3 
HC1O 4 

Not all elements form a complete series of acids, but the 
nomenclature usually agrees with the above principles. 

92 Descriptive Chemistry. 

Some acids have commercial names. Thus, sulphuric acid 
is often called oil of vitriol, and hydrochloric acid is known 
as muriatic acid. Acids in which carbon is the essential 
component end hi -ic, but they are often arbitrarily named 
(see Organic Acids). 

An examination of the formulas of acids shows that all do not con- 
tain the same number of hydrogen atoms. Acids are sometimes classi- 
fied by the number of hydrogen atoms which can be replaced by a metal. 
This varying power of replaceability is called basicity. A monobasic 
acid contains only one atom of replaceable hydrogen in a molecule, e.g. 
nitric acid, HNO 3 . A molecule of acetic acid (C 2 H 4 O 2 ) contains four 
atoms of hydrogen, but for reasons which are too complex to state here, 
only one of these atoms can be replaced by a metal. Dibasic and 
tribasic acids contain two and three replaceable hydrogen atoms, e.g. 
sulphuric acid (H 2 SO 4 ) and phosphoric acid (H 3 PO 4 ). Obviously, 
monobasic acids form only one class of salts, dibasic acids form two 
classes, tribasic acids form three, and so on. 

Bases. The term base, in a narrow sense, means the 
strong bases, which are very soluble in water, and are com- 
monly known as alkalies, e.g. sodium, potassium, and 
ammonium hydroxides. In a broad sense it means any 
substance which will neutralize an acid, e.jr. calcium oxide, 
ammonia gas, as well as the hydroxides of metals. Most 
bases are solids ; but since they are usually soluble in water, 
these solutions, as in the case of acids, are familiarly 
called the base, or alkali, itself. Concentrated alkalies, 
like concentrated acids, are corrosive. The common alka- 
lies sodium and potassium hydroxides are often called 
caustic soda and caustic potash to emphasize this property ; 
and calcium oxide, or lime, is sometimes called caustic 
lime ; the corrosive nature of ammonium hydroxide, or 
ordinary 'ammonia, is also well known. Substances which 
turn red litmus to blue are said to contain an alkali (or 
base), to be alkaline, or to have an alkaline reaction. 

Acids, Bases, and Salts. 93 

The word basic is often used instead of alkaline. Other 
tests besides that with litmus must be applied, however, to 
determine the exact nature of a substance having an alka- 
line reaction. Alkalies dissolve grease and fats, and are 
often used as cleansing agents, ammonium hydroxide 
being widely employed for this purpose. They also inter- 
act with fats to form soaps, large quantities of sodium 
hydroxide being annually utilized in the soap industry (see 

A base, like an acid, is rather difficult to define. We 
might say that a base is an hydroxide or oxide of a metal, 
which will neutralize an acid, thereby forming a salt. 
The term must include ammonia, which does not contain 
a metal. But, as we shall see later, a certain combination 
of elements related to ammonia acts like a metal (see 

Nomenclature of Bases. There is no general rule 
covering the nomenclature of bases, as in the case of 
acids. Since most bases contain hydrogen and oxygen, 
they are often called hydroxides. Hydrate is sometimes 
used as a synonym of hydroxide. The term alkali em- 
phasizes general properties rather than suggests specific 
composition. Hydroxides are distinguished from each 
other by placing the name of the metal before the word 
hydroxide, e.g. sodium hydroxide, potassium hydroxide, 
calcium hydroxide. The common hydroxides have long 
been known by several names. Thus, calcium hydroxide 
is often called limewater. Ammonium hydroxide is some- 
times called ammonia water or simply (but inaccurately) 
ammonia, and it was formerly called volatile alkali. Be- 
sides the common names of the hydroxides of sodium and 
potassium already given, they are sometimes called fixed 

94 Descriptive Chemistry. 

Not all bases contain the same number of hydroxyl groups. Hence 
bases, like acids, may form one or more salts. This power is called 
acidity. Bases are called monacid, diacid, triacid bases, etc., accord- 
ing to the number of replaceable hydroxyl groups present in a molecule. 
Thus, calcium hydroxide (Ca(OH) 2 ) is a diacid base, and aluminium 
hydroxide (A1(OH) 3 ) is a triacid base. 

Salts. Sodium chloride, or ordinary table salt, is the 
most familiar salt. It has been known for ages. Doubt- 
less this class of chemical compounds received its name 
because of the general resemblance most of them bear to 
common salt. Most salts are solid and are soluble in 
water. Many of them have no action on litmus, and are, 
therefore, said to be neutral or to have a neutral reaction. 
This indifference to litmus is not a decisive test for a 
salt, since many other substances, water for example, 
have no action on litmus. Nevertheless the term neutral 
is applied to substances which do not change the color 
of litmus. 

Some substances which are salts, as far as their structure and method 
of formation are concerned, do not have a neutral reaction. Thus, 
sodium carbonate, which is the sodium salt of carbonic acid, has a 
marked alkaline reaction, being in fact known in commerce simply as 
" alkali." 

A salt may be defined as the main product of the inter- 
action of an acid and a base. It may, however, be a sub- 
stance which has the properties of a salt, regardless of the 
method of formation. 

Salts are formed in various ways. The interaction of an acid and a 
base has been mentioned. The interaction of acids with oxides of cer- 
tain metals or with metals themselves produces salts. Sodium oxide 
and sulphuric acid interact and form the salt sodium sulphate, thus : 

Na,O + H 2 SO 4 Na 2 S0 4 + H 2 O 

Sodium Oxide Sulphuric Acid Sodium Sulphate Water 

Acids, Bases, and Salts. 95 

While zinc and sulphuric acid, as already stated, form the salt zinc 
sulphate as well as hydrogen, thus : 

Zn + H 2 S0 4 = ZnS0 4 + H 2 

Zinc Sulphuric Acid Zinc Sulphate Hydrogen 

Carbonates interact with acids and form other salts. Calcium carbonate 
and hydrochloric acid form the salt calcium chloride, thus : 

CaC0 3 + 2HC1 = CaCl 2 + CO 2 + H 2 O 
Calcium Hydrochloric Calcium Carbon Water 
Carbonate Acid Chloride Dioxide 

Nomenclature of Salts. The name of salts containing 
oxygen are derived from the name of the corresponding 
acid. The characteristic suffix of the acid is changed to 
indicate this relation. Thus, the suffix -ic becomes -ate, 
and the suffix -ous, becomes -ite. Hence : 

Sulphuric acid forms sulphates. 
Sulphurous acid forms sulphites. 
Nitric acid forms nitrates. 
Nitrous acid forms nitrites. 
Chloric acid forms chlorates. 
Hypochlorous acid forms hypochlorites. 
Permanganic acid forms permanganates. $ 

The name of the replacing metal is retained, e.g. potas- 
sium chlorate, sodium sulphate, calcium hypochlorite, po- 
tassium permanganate. Notice that the prefixes hypo- and 
per- are not changed. 

The names of salts containing only two elements, fol- 
lowing the general rule for binary compounds, end in -ide. 
This suffix is added to a modification of the name of the 
non-metal, giving the names chloride, bromide, sulphide, 
fluoride, etc. The prefix hydro- which is contained in the 

96 Descriptive Chemistry. 

name of the acid is omitted. Thus, the name of the 
sodium salt of hydrochloric acid is sodium chloride ; simi- 
larly, there are the names potassium chloride, calcium 
fluoride, and sodium iodide. Sometimes, the salts of these 
hydrogen acids are called halides to emphasize their rela- 
tion to common salt, which in Greek is called halos. 

Salts in which all the hydrogen atoms of the corresponding acid 
have been replaced by a metal are called normal salts, e.g. sodium 
sulphate, Na.,SO 4 . If some of the hydrogen atoms are not replaced by 
a metal, an acid salt is formed. Thus, acid sodium sulphate may be 
regarded as derived from sulphuric acid, which is dibasic, by replacing 
one of the atoms of hydrogen by sodium, though of course the salt is 
not prepared in this way. Expressed as formulas these relations may 
be written thus : 

Acid Acid Salt Normal Salt 

H,SO 4 HNaSO 4 Na 2 SO 4 

Only those acids which contain two or more replaceable hydrogen 
atoms form acid salts. On the other hand, if not all the hydroxyl 
groups of a base are replaced when the base reacts with an acid, then a 
basic salt results. Thus, basic nitrate of bismuth may be regarded as 
the salt derived from bismuth hydroxide (Bi(OH) 3 ) by replacing one 
hydroxyl group of the base by the group NO 3 of nitric acid. The 
formula of this basic nitrate of bismuth is Bi(OH) 2 NO 3 . 

The following equation illustrates the changes : 

Bi(OH) 3 + HN0 3 = Bi(OH) 2 NOo + H 2 O 

Bismuth Hydroxide Nitric Acid Basic Bismuth Nitrate Water 

Only those bases having two or more hydroxyl groups can form basic 
salts. Some basic salts are very complex. 

Relation of Oxides to Acids and Bases. Most non- 
metallic elements form oxides which unite with water and 
produce an acid. The oxides of many metallic elements, 

Acids, Bases, and Salts. 97 

on the other hand, unite with water and produce hydrox- 
ides. The two oxides of the non-metal sulphur act thus 

50 2 + H 2 O = H 2 SO 3 
Sulphur dioxide Water Sulphurous Acid 

50 3 + H 2 O = H 2 SO 4 
Sulphur Trioxide Water Sulphuric Acid 

The oxide of the metal calcium acts thus 

CaO -f H 2 O = Ca(OH) 2 
Calcium Oxide Water Calcium Hydroxide 

Oxides of non-metals which unite with water and thereby 
produce acids are called anhydrides, i.e. literally, sub- 
stances without water. Examples are carbonic anhydride 
(CO 2 ), sulphuric anhydride (SO 3 ), phosphoric anhydride 
(P 2 O 5 ). Oxides of metals which produce hydroxides are 
called basic oxides. A few oxides behave exceptionally. 
It is convenient to regard an anhydride as the root or 
basis of its corresponding acid, and a basic oxide as the 
root of its hydroxide. 

The fact that many non-metallic oxides redden moist blue litmus led 
Lavoisier into the erroneous belief that oxygen is an essential compo- 
nent of acids. And some authorities even now (incorrectly) speak of 
these oxides as acids ; thus, carbon dioxide (CO 2 ) is occasionally called 
carbonic acid. The compounds which Lavoisier galled acids were anhy- 
drides. And it was not until about 181 1 that Davy showed (i) that some 
acids do not contain oxygen (e.g. hydrochloric acid, HC1 ), and (2) that 
the so-called acids of Lavoisier are not real acids until they have obtained 
hydrogen from the water with which they combine. 

Neutralization has been defined as the series of changes 
whereby acids and bases mutually destroy each other's 
characteristic properties and produce a salt and water. 


Descriptive Chemistry. 

But neutralization has a deeper meaning and broader ap- 
plication than the mere destruction 
of properties. 

If measured volumes of different acids 
are exactly neutralized by different alkalies, 
remarkable relations are revealed. This may 
be done by dropping one into the other from 
a graduated tube, called a burette (Fig. 15). 
The exact point of neutralization is shown 
by an indicator; this is a solution of litmus 
or some other substance, which tells by the 
color whether the solution is acid or alkaline. 
Experiment shows that (i) a definite quan- 
tity of an acid neutralizes a definite quantity 
of an alkali, (2) the same acid is neutralized 
by different quantities of different alkalies, 
and (3) the ratio of the quantities of the 
FIG. 15. Burettes. different alkalies is the same for all acids. 1 


1. Define and illustrate (a) an acid, (6) a base, (c) a salt, (d} an al- 
kali, (e) hydroxyl, (/) an hydroxide. 

2. Name three common acids and bases. State the general proper- 
ties of each class. 

3. Define and illustrate (a) neutralization, (b} acidity of bases, (c} 
basicity of acids, (</) normal, acid, and basic salts, (V) caustic alkali, 
(/) radical. 

4. What is the literal meaning of (a) acid ( adj.), () caustic, (c) per-, 
(d) hypo-, (i) anhydride? 

5. Name the sodium salt of hydrochloric acid. Name the corre- 
sponding salt of potassium, lead, calcium, barium, zinc, silver. 

6. Name the same salts of nitric acid. Of nitrous acid. 

7. Name the same salts of sulphuric acid. Of hypochlorous acid. 
Of perchloric acid. 

1 A more extended treatment of this subject may be found in the author's 
"Experimental Chemistry," pp. 124 ff. 

Acids, Bases, and Salts. 99 

8. Name the hydroxides corresponding to sodium, potassium, calcium, 
barium, zinc, lead, copper. 

9. Name the potassium salt of manganic acid, calcium salt of hydro- 
fluoric acid, sodium salt of carbonic acid, potassium salt of tartaric acid, 
lead salt of chromic acid, potassium salt of hydrobromic acid, potassium 
salt of permanganic acid. 


Review any of the preceding problems, especially those in Appendix, 



Equivalents. The equivalent or equivalent weight of 
an element is that weight which is chemically equivalent 
to one part by weight of hydrogen. More specifically, it 
is the number of grams of an element which liberates, 
replaces, or combines with I gm. of hydrogen. Ex- 
periments show that approximately 32.5 gm. of zinc will 
liberate I gm. of hydrogen from an acid. Hence 32.5 
is the equivalent of zinc. Similarly, 23 gm. of sodium 
liberate I gm. of hydrogen from water. A summary of 
numerous experiments reveals the following 






(by definition) 









3 2 -5 






2 3 








Atomic ' Weights. \ "- *." ' ; \ /, \\ \ : /, i o i 

Analysis of chemical compounds determines the propor- 
tion of their components by weight. And in many cases 
such experiments verify the equivalents found by other 
methods. Thus, experiment shows that 

35-5 P ar *s of chlorine unite with 23 of sodium, or 39 of potassium. 

80 parts of bromine unite with 23 of sodium, or 39 of potassium. 

108 parts of silver replace 23 of sodium, or 39 of potassium. 

The above elements always unite in these proportions. 
But some elements unite in several proportions. Thus, 
eight parts by weight of oxygen combine with one part of 
hydrogen to form water. But in a large number of com- 
pounds sixteen parts of oxygen combine with various parts 
of different elements. Similarly, nitrogen unites in the 
proportion of fourteen, twenty-eight, and forty-two parts 
by weight with different parts of other elements. In a 
word, there are multiples of equivalents. Comparison 
shows a striking coincidence between many equivalent 
weights and the accepted atomic weights of the same 
elements. This topic is discussed and applied in Chapter 

Atomic Weights. One of the essential properties of 
matter is weight. According to the atomic theory, atoms 
have weight. But the weight of an atom is so small that 
we cannot determine it. We can, however, find the rela- 
tive weight of an atom ; that is, how many times heavier 
one atom is than another atom. If we adopt one as the 
weight of an atom of hydrogen, the weights of atoms of 
other elements can be readily expressed in terms of this 
standard. Thus, when we say the atomic weight of sodium 
is twenty-three, we mean that an atom of sodium weighs 
twenty-three times as much as an atom of hydrogen. The 

IO2 Descriptive Chemistry. 

determination of the exact atomic weight of an element is 
a difficult task. Many principles influence the final selec- 
tion of the number adopted as the atomic weight. We 
have already seen that there is a definite relation between 
the equivalent weight and the atomic weight of an element. 
But this method cannot be used exclusively to determine 
atomic weights, because it does not enable us to tell the 
number of atoms in a molecule. There is also a definite 
relation between the molecular weight of a compound and 
the atomic weights of the elements in the compound. 
These topics and others related to them will be discussed 
in Chapter XIII. For the present, the approximate atomic 
weights found in the Appendix, 5, may be used in solv- 
ing problems and interpreting equations. 

The atomic weights are not necessarily whole numbers, but they are 
nearly so in many cases, and for most purposes round numbers may be 
used. Different atomic weights are sometimes given for the same ele- 
ment. This is due (i) to the disagreement among chemists as to the 
accuracy of certain results, and (2) to the use of several standards for 
reckoning atomic weights. For many years hydrogen was the standard. 
But for scientific reasons oxygen is being adopted as the standard, and 
1 6 is accepted as its atomic weight. This change does not alter the 
facts; it merely changes the relative values of the atomic weights. 
Thus, the atomic weight of hydrogen becomes 1.008, if oxygen equals 
1 6, and others are proportionally changed. 

Tables of atomic weights have been prepared on both standards 
(H = i and O = 16). Both tables are given in the Appendix, 5. 

Symbols and Atomic Weights. Symbols not only rep- 
resent atoms, but they express atomic weights. Thus, O 
represents one atom of oxygen, but it also means that this 
atom weighs sixteen times more than an atom of hydro- 
gen. Similarly, K represents an atom of potassium, 
which weighs thirty-nine times more than an atom of 

Chemical Calculations. 103 

Molecular Weights. Since atoms combine to form 
molecules, a molecular weight is the sum of the weights 
of the atoms in a molecule. A molecule of nitric acid 
contains one atom each of hydrogen and nitrogen, and 
three atoms of oxygen ; hence its molecular weight is 
i + 14+ 16 x 3 = 63. Given the formula, the molecular 
weight is easily found by adding the atomic weights. 
The molecular weight and formula of a compound, there- 
fore, are rigidly connected; and just as a symbol stands 
for an atomic weight, so a formula expresses a molecular 
weight. It is customary to assume the simplest formula 
(i.e. the one corresponding to the lowest molecular weight) 
until experiments show which is the correct one. 

Many facts and principles determine the final selection of the molecular 
weight, and hence the formula, of a compound. These will be discussed 
in Chapter XIII. 

Chemical Calculations are largely based on atomic and 
molecular weights. 

Percentage Composition. Since the formula of a com- 
pound expresses its composition, it is possible to calculate 
from the formula the composition in per cent. The for- 
mula of sulphuric acid is H 2 SO 4 , and its molecular weight 
is 98, i.e. 2 + 32+64. The calculations are most easily 
made by the following proportions: 

2 : 98 : : x : 100, x = 2.04 per cent of hydrogen. 
32 : 98 : : x : 100, x 32.65 per cent of sulphur. 
64 : 98 : : x : 100, x = 65.31 per cent of oxygen. 
Total 100.00 per cent. 

By the same method the percentage composition of any 
compound may be calculated. 

IO4 Descriptive Chemistry. 

Simplest Formula. The simplest formula of a compound 
may be found by dividing the percentage of each element 
in the compound by its atomic weight. The percentage 
composition of sulphuric acid is H = 2.04, 8 = 32.65, 
= 65.31. Dividing each percentage by the atomic 
weight of the element, we have(approximately)2.O4 ^-1=2, 
32.65-^-32=1, 65.31-^16 = 4. Hence the simplest for- 
mula of sulphuric acid is H 2 SO 4 . Sometimes the prod- 
ucts of the percentages divided by the atomic weights 
are not whole numbers. In that case the simplest relation 
is found by proportion. The following problem illustrates 
this principle : the percentage composition of a compound 
is C = 40, H = 6.67, = 53.33. Dividing as above, we 
have 40 H- 12 = 3.33, H-^ i =6.67, 53-33^ 16=3.33. But 
3.33, 6.67, 3.33 are in the same proportion as 1:2:1. 
Hence the simplest formula is CH 2 O. 

Quantitative Significance of Equations. It is possible 
to express reactions in the form of equations because in 
every chemical change no weight is lost or gained. 

It has already been stated that the equation for the 
reaction between magnesium and oxygen is 

Mg + O MgO 

Magnesium Oxygen Magnesium Oxide 

This equation is the outcome of the following : it can be 
readily shown by experiment that when magnesium is 
heated in air or oxygen, the magnesium and oxygen com- 
bine in the ratio 3 : 2. Now results like this are usually 
expressed in terms of the atomic weights of the reacting 
elements. But we do not know the number of atomic 
weights of these elements which must be taken to produce 
the ratio 3 : 2. That is, we do not know whether the 
ratio requires the atomic weight or some multiple of it. 

Quantitative Significance of Equations. 105 

But, if we let y equal the unknown number of atomic 
weights of magnesium and z the unknown number of 
atomic weights of oxygen, then we can write the prelimi- 
nary equation thus 

y X at. wt. of mag. : z X at. wt. oxygen = 3:2. 

The atomic weight of magnesium is 24 and of oxygen is 
1 6. Therefore the problem reduces itself to finding the 
values of y and z in the equation 

y x 24 : s x 16 = 3: 2. 

Obviously, y = i and z = i. Now the symbol Mg stands 
for 24 parts of magnesium and O for 16 parts of oxygen. 
That is, Mg not only means one atom of magnesium, but 
also that this atom weighs 24, if one atom of oxygen 
weighs 1 6. Therefore, Mg and O represent the number 
of atoms which are equivalent arithmetically to the ratio 
3 : 2 found by experiment. Since one atom of magne- 
sium and one of oxygen unite to form magnesium oxide, 
its formula is MgO. Therefore, the final equation is 

Mg-f O=MgO. 

Again, suppose we wish to find the correct equation for 
the reaction between hydrogen and oxygen in the forma- 
tion of water. Experiment shows that hydrogen and oxy- 
gen combine in the ratio of i : 8 by weight. Pursuing 
the same line of argument as above, we let y equal the 
unknown number of atomic weights of hydrogen, and z 
that of oxygen. The preliminary equation is 

y x at. wt. of hydrogen : z x at. wt. of oxygen = i : 8. 
The atomic weight of hydrogen is i and of oxygen is 16. 
The equation now becomes 

y x i : z x 16 = i : 8. 

106 Descriptive Chemistry. 

Obviously, y = 2 and z = i. Now the symbol H stands 
for i part of hydrogen and O for 16 parts of oxygen. 
Therefore, 2 H and O represent the number of atoms which 
corresponds to the ratio i : 8, found by experiment. Since 
two atoms of hydrogen and one atom of oxygen unite to 
form water, its formula must be H 2 O. Therefore the final 
equation is- H 2 + O = H 2 O. 

By a similar treatment, the experimental foundation of 
all equations can be shown. 

Equations illustrating Reactions. The simplest equation for the 
preparation of oxygen from mercuric oxide is 

HgO = Hg + O / 

Mercuric Oxide Mercury Oxygen 

When sulphur and carbon are burned in air or in oxygen, the equations 

are ~ s + o 2 = so, 

Sulphur Oxygen Sulphur Dioxide 

C + O, = C0 2 
Carbon Oxygen Carbon Dioxide 

The equation for the preparation of hydrogen from zinc and hydro- 
chloric acid is 

Zn + 2 HC1 = H, + ZnCl, y 

Zinc Hydrochloric Acid Hydrogen Zinc Chloride 

When hydrogen burns, the equation is 

H, + O = H a O 
Hydrogen Oxygen Water 

The equation for the formation of water is the same, though it is some- 
times written- 2 H 2 + O 2 = 2 H,O. 

The equation for the reaction in determining the gravimetric composi- 
tion of water is 

CuO + H 2 = H 2 O + Cu Y 

Copper Oxide Hydrogen Water Copper 

Problems based on Equations. 107 

The interaction of sodium and water is represented thus 

Na + H 2 H + NaOH 

Sodium Water Hydrogen Sodium Hydroxide 

When phosphorus burns in air (or oxygen), the simplest equation is 

2? + 50 - PA 

Phosphorus Oxygen Phosphorus Pentoxide 

Problems based on Equations. Since equations are expressions 
of chemical reactions which involve no loss in weight, it is possible to 
solve many problems connected with reactions. An equation states 
the proportions which participate in a reaction. Obviously, any con- 
venient weights of zinc and sulphuric acid might be allowed to interact, 
but the factors and products are always in the proportions given in the 

Zn + H 2 SO 4 = H 2 + ZnSO 4 
Zinc Sulphuric Acid Hydrogen Zinc Sulphate 

65 98 2 161 

This expression means that 65 parts of zinc always interact with 98 
parts of sulphuric acid and yield 2 parts of hydrogen and 161 parts of 
zinc sulphate. For parts we may read grams, ounces, kilograms, any 
unit, but the same unit must be used throughout the calculations. 
Therefore, if we know the weight of one substance participating in a 
reaction, all other weights involved may be readily calculated. 

Suppose 45 gm. of zinc interact with sulphuric acid ; the weights 
of (a) acid required, (b) hydrogen formed, and (c) zinc sulphate produced 
are found by the following proportions : 

(a) 65 : 98 { : 45 : y, x= 67.8 gm. sulphuric acid. 
() 65 : 2 : : 45 : x, ^=1.38 gm. hydrogen. 
(c} 65 : 161 : : 45 : x, x= II 1.4 gm. zinc sulphate. 

Hence to solve similar problems, first write the equation with the correct 
atomic or molecular weights, 1 and then state the problem in the form 
of a proportion like those given above. 

i The atomic weights are given in the table in the Appendix, 5. Molecular 
weights are obtained by adding the proper atomic weights. 

io8 Descriptive Chemistry. 


1 . Define and illustrate the term equivalent. What is the equivalent 
of hydrogen, oxygen, sulphur, zinc, copper, magnesium, silver, potassium, 
aluminium ? 

2. What is the equivalent of chlorine and of bromine? 

3. How are equivalents determined? Are they the result of theory 
or actual analysis ? 

4. Expand the topic, " Atomic weights are often multiples of equiva- 

5. What is the atomic weight of an element? "How is it related to 
the equivalent weight of the element? Is an atomic weight absolute or 
relative ? What is the standard of atomic weight? 

6. What does O represent besides one atom of oxygen? 

7. What is the approximate atomic weight of hydrogen, oxygen, 
and sodium? 

8. What is meant by molecular weight? Illustrate by nitric acid or 
potassium chlorate. What is the relation between formula and molecular 
weight ? 

9. Define and illustrate (a) percentage composition, and () sim- 
plest formula. 

10. How do we know that the correct equation for the combination 
of magnesium and oxygen is Mg + O = MgO ? That S + O 2 = SO 2 is the 
correct equation for the combination of sulphur and oxygen ? 


1 . Calculate the percentage composition of (a) water (H 2 O), (b} zinc 
sulphide (ZnS), (c} zinc carbonate (ZnCO 3 ), (d) potassium chlorate 
(KC10 3 ). 

2. Calculate the percentage composition of (#) sugar (C 12 H 22 O 11 ), 
() calcium sulphate (CaSO 4 ), (c) zinc sulphate (ZnSO 4 ), (d) magne- 
sium oxide (MgO), (e) copper oxide (CuO). 

3. Calculate the molecular weight of the following compounds by 
finding the sum of the atomic weights : (a) copper sulphate (CuSO 4 ), 
(If) barium chloride (BaCL,), (c) manganese dioxide, (d} calcium oxide, 
(e) sodium hydroxide, (/) potassium hydroxide, (g) sodium carbonate, 
(^) potassium nitrate (KNO 3 ), 

Problems. 109 

4. Calculate the simplest formula of the compounds which have the 
indicated composition, and give the name of each compound : (a) 
H = II. 11,0 = 88.89; ()Na = 32.39,0 = 45-o7,S = 22.54; (Y) = 27.27, 
= 72.72. 

5. Calculate the simplest formula of the compounds which have the 
following composition : (a) N = 82.353, H = 17.647 ; () O = 30, Fe = 70 ; 
(c) H = i, C = 11.99, O =47-95> K = 39.06. 

6. How much oxygen can be prepared from (ft) 122.5 m< f potas- 
sium chlorate, (b) 245 gm., and (c) 421 gm. ? 

Solution. The equation is 

KC10 3 = 3 + KC1 
122.5 =48 4- 74-5 

These equation weights are obtained by adding the atomic weights found 
in the table, (a) By inspection, 122.5 g m - of potassium chlorate yield 
48 gm. of oxygen. (b) The proportion needed is 122.5 : 48 :: 245 : x. 
And x = 96 gm. (c) Similarly, 122.5 : 4^ : : 4 21 ' * And x 164.9. 

7. (a) How much oxygen can be prepared from 50 gm. of potassium 
chlorate, and (b) how much potassium chloride will remain? 

Ans. (a} = 19.59, (b) = 30.41. 

8. A certain weight of potassium chlorate was heated until completely 
decomposed. The residue weighed 20.246 gm. (a) What was its weight ? 
(b) How much oxygen was evolved? Ans. (a)= 33.29, (b) 13.044. 

9. What weight of potassium chlorate is needed to generate 144 gm. 
of oxygen? .^^.367.5. 

10. What weight of potassium chloride remains after obtaining 8 gm. 
of oxygen from potassium chlorate? Am. 12.416. 

1 1 . How many grams of oxygen can be generated from 490 gm. of po- 
tassium chlorate? Ans. 192. 

12. How much hydrogen can be prepared from (a) 65 gm. of zinc, 
(b) 130 gm., (V) 297 gm.? Ans. (c} 9.14- 

13. How much zinc is needed to prepare (a) 2 gm. of hydrogen, (b) 
14 gm., and (c) 17 gm.? 

14. How much zinc sulphate can be prepared from (a) 98 gm. of sul- 
phuric acid, (b) 196 gm., and (c) 427 gm.? Ans. (c) 701.5. 

15. A balloon holds 132.74 kg. of hydrogen. How much (a) zinc 
and (b) sulphuric acid are needed to produce the gas? 

Ans. (a) 43 I 4-05> (*) 6504.26. 

no Descriptive Chemistry. 

1 6. How much (a) mercury and () oxygen can be obtained from 
10 gm. of mercuric oxide? (Equation is HgO = Hg + O, or 216 = 
200 + 1 6.) Ans. (a) 9.259, (b} 0.74. 

17. How much mercury will remain after obtaining 48 gm. of oxygen 
by heating mercuric oxide? 

1 8. A lump of carbon weighing 24 gm. is burned in air. What weight 
of (a) carbon dioxide is formed and () oxygen is needed? (c) If a liter 
of oxygen weighs 1 .43 gm., what volume of oxygen is needed ? (Equation 
is C + O 2 = CO 2 , or 12 + 32 = 44.) Ans. (c) 44.75 1. 

19. What weight of carbon dioxide is formed by burning 112 Ib. of 
coal containing 15 per cent of impurities? 

20. A lump of sulphur weighing 32 gm. is burned in air. Calculate 
the weight of (a) oxygen needed and (&) sulphur dioxide formed. 
(Equation is S + O 2 = SO 2 , or 32 + 32 = 64.) 

21. Calculate the weight of oxygen needed to burn 731 gin. of sul- 
phur containing 15 per cent of impurities. Ans. 621.35. 

22. What weight of sulphur dioxide is formed by burning 67 per cent 
of 8794 kg. of sulphur? 


CHEMICAL action is always manifested by one or more 
of the different forms of energy, such as light, heat, and 
electricity. This means that a chemical change involves 
not only a rearrangement of matter, but also a transfor- 
mation of energy. Thus, when coal is burned, a new 
compound called carbon dioxide is formed, but heat is also 
liberated. Sometimes we pay more attention to the result- 
ing matter than to the energy, but both are involved. In 
the present chapter we shall emphasize the relation of 
energy to chemical action. The law of the conservation 
of energy should be recalled in this connection. Energy, 
like matter, cannot be created or destroyed ; we can only 
transform it. And the transformation involves no loss or 
gain. Hence, chemical energy, which is the immediate 
cause of chemical action, will appear as heat, light, or 

The Relation of Light to Chemical Action is illustrated in photogra- 
phy. Coatings consisting of compounds of silver and organic matter are 
quickly blackened by light (see Photography). Sunlight fades many 
colors. It likewise assists the chemical changes involved in the growth 
of plants. The formation of the green coloring matter of foliage is 
partly due to sunlight. A mixture of hydrogen and chlorine gases re- 
mains unchanged in the dark, but in direct sunlight it explodes violently. 
On the other hand, light is often a product of chemical action. Many 
chemical experiments show this, especially those with oxygen. Sparks, 
most flames, and the flash of a gun are other illustrations of the close 
relation between light and chemical action. Combustion in its varied 
forms is also manifested by light, as well as by heat. 


H2 Descriptive Chemistry. 


Heat and Chemical Action are closely and definitely 
related. Every chemical change is attended by the libera- 
tion or absorption of heat. Moreover, the heat involved 
can often be measured. Heat is measured in calories, a 
calorie being the quantity of heat necessary to raise the 
temperature of I gm. of water from o to i C. For 
example, the heat liberated by the burning of I gm. of 
hydrogen is 34,200 cal., and of I gm. of pure charcoal is 
about 8000 cal. Attention has already been called to the 
high temperature of the hydrogen flame (see Chapter III). 

Ordinary chemical equations do not express changes in energy. To 
represent heat changes, the number of calories of heat involved is 
placed after the equation, thus : 

H 2 + O = H 2 O + 68,400 cal. 
Hydrogen Oxygen Water 

This is called a thermal equation, and it means that 68,400 cal. of heat 
are liberated, when 2 gm. of hydrogen unite with 16 gm. of oxygen to 
form 18 gm. of water. In some changes heat disappears. Thus, when 
carbon unites with sulphur to form carbon disulphide, heat is absorbed. 
The equation expressing this fact is 

C + S 2 = CS 2 19,600 cal. 

Carbon Sulphur Carbon Disulphide 

Heat involved in the formation of a particular compound is called heat 
of formation of that compound. If heat is liberated in the formation 
of a compound, the heat is called positive ( + ); and the compound is 
termed exothermic. Heat of formation which is absorbed is called 
negative ( ) ; and a compound having a negative heat of formation is 
said to be endothermic. Exothermic compounds are stable, and can 
be decomposed only by the addition of the same quantity of heat liber- 
ated by their formation. Thus, 68,400 cal. of heat, or an equivalent 
quantity of energy, must be added to 18 gm. of water to decompose it 

Light, Heat, Electricity, and Chemical Action. 113 

into 2 gm. of hydrogen and 16 gm. of oxygen. Such heat is called 
heat of decomposition. On the other hand, endothermic compounds 
are unstable, and often explosive. They decompose easily with the 
liberation of heat. Ozone is endothermic. Heat is absorbed during 
its formation from oxygen ; but when ozone decomposes, heat is liber- 
ated. Two parts (by volume) of ozone form three parts (by volume) 
of oxygen and liberate 72,400 cal. 

A familiar instance of the evolution of heat by chemical 
action is the slaking of lime. When lime and water are 
mixed, their union produces sufficient heat to boil water 
and often to set fire to wood. Steam can be seen escaping 
from the boxes in which lime is being mixed with water and 
sand to form plaster or mortar. Buildings in which lime is 
stored sometimes take fire, if rain leaks in upon the lime. 
Ships loaded with lime are in constant danger of being 
burned. Other substances liberate heat when added to 
water, e.g. sulphuric acid, sodium and potassium hydrox- 
ides, and the metals, sodium and potassium. 

Heat is the initial cause of many chemical changes. It 
is necessary to start many reactions, just as a stone on top 
of a hill must be pushed before it will roll toward the 
bottom. Hydrogen and oxygen mix freely without com- 
bining, but union occurs the instant heat is applied in form 
of a flame or an electric spark. Similarly, illuminating 
gas must be lighted, i.e. raised to the kindling tempera- 
ture before the chemical changes which cause the light 
and heat can proceed. These facts mean that chemical 
action often depends upon temperature. This statement 
has been strikingly illustrated in the last four years. At the 
extremely low temperature obtained by using liquid air and 
similar substances, it appears that many chemical reactions 
cease. While at the exceedingly high temperature pro- 
duced by electricity many changes, chemical and physical, 
hitherto impossible, occur quickly and simply. 

114 Descriptive Chemistry. 

The Electric Furnace of Moissan. Until recently the 
heat needed for chemical changes was obtained by burn- 
ing carbon or its compounds, such as charcoal, illumi- 
nating gas, and oil. Sometimes the blast lamp and 
oxyhydrogen blowpipe were used. But all these sources 
have been surpassed in efficiency by the electric furnace. 

It is well known that an electric arc light produces in- 
tense heat. The high temperature of the arc, i.e. space 
between the glowing ends of the carbons, is unequaled by 
that of any other source of artificial heat. If the carbon 
rods are inclosed in a box that prevents the escape of heat, 
a temperature estimated to be about 3500 C. is produced 
inside the box. This apparatus is called an electric furnace. 
It was devised and perfected by the French chemist, Mois- 
san, and used by him in experimenting at high temperatures. 
One form of the electric furnace is shown in Figure 16. 

FIG. 16. Moissan's electric furnace. 

Moissan's description of this furnace is as follows : " It consisted of 
two bricks of quicklime placed one on top of the other. The lower 
brick contained a longitudinal groove to receive the two electrodes 
[carbon rods], and situated in the center was a small cavity. This 
cavity might vary in size, and contained a bed some centimeters in 
depth of the substance to be acted upon by the heat of the arc, or a 
small crucible of carbon containing the substance to be treated may be 
placed there. The upper brick was slightly hollowed out in the part 
just above the arc. As the intense heat of the current soon melted the 




Light, Heat, Electricity, and Chemical Action. 115 

surface of the lime, giving it, at the same time, a beautiful polish, a 
dome was obtained in this way which reflected all the heat on to the 
small cavity which contained the crucible." Figure 17 is a vertical sec- 
tion of the furnace, showing the parts slightly separated. The furnace 
is small, some being only 16 to 18 cm. (about 7 in.) long, 15 cm. wide, 
and 14 cm. high. 
The carbon rods are 
from i to 5 cm. in 

When a cur- 
rent is passed 

through the Car- FIG. 17. Vertical section of Moissan's electric furnace. 

bon rods, the 

tremendous heat produced is retained in the space by the 
non-conducting walls and acts upon the substance below 
the arc. The outside of the furnace remains cold enough 
to be touched by the hand, but the inside is almost twice 
as hot as the oxyhydrogen flame. There is no electrical 
action upon the chemicals. The intense heat alone pro- 
duces the remarkable changes, which are often accom- 
plished in a few minutes. Sand, lime, magnesium oxide, 
and other refractory oxides melt and volatilize. The ele- 
ments carbon, silicon, and boron boil; and gold, copper, 
and platinum quickly melt and vaporize. Large masses 
of rare and uncommon elements are quickly reduced from 
their oxides and obtained in the pure state, e.g. chromium, 
manganese, tungsten, uranium, and molybdenum. Char- 
coal becomes graphite. And stable compounds of carbon, 
boron, and silicon are formed. These are the carbides, 
borides, and silicides. Some of the carbides have an in- 
dustrial use as well as scientific interest, especially calcium 
carbide and silicon carbide (see below). Other carbides 
are the sources of pure metals, since the fusion of a car- 
bide and oxide of the same metal yields the metal itself. 

ii6 Descriptive Chemistry. 

Industrial Use of the Electric Furnace. Huge elec- 
tric furnaces constructed on the type devised by Moissan 
are in active operation. And since electricity is now ob- 
tained in many localities by running dynamos by water, 
new industries requiring intense and continuous heat have 
recently sprung into existence. Several of these plants 
are located at Niagara Falls, which furnishes enormous 
power at a relatively small expense. 

Calcium Carbide is made on a large scale by heating a 
mixture of lime and coke (a form of carbon) in an electric 
furnace. The chemical change is caused solely by the 
intense heat and may be represented thus : 

3C + CaO = CaC 2 + CO 
Carbon Lime Calcium Carbide Carbon Monoxide 

This method of making calcium carbide cheaply was dis- 
covered independently and at about the same time (1892- 
1895) by Moissan and Willson. The furnaces now in 
operation vary in details, but all have one essential feature, 
viz., the heat is generated by an electric current passing 
between two carbon electrodes. In most furnaces one elec- 
trode is a crucible wholly or partly of carbon, and the other 
electrode is a stout carbon pillar dipping into the mixture. 
Calcium carbide is a hard, brittle, dark gray, crystalline 
solid with a metallic luster. Its specific gravity is 2.2. 
The most striking and useful property is its action with 
water, acetylene being formed, thus: 

CaC 2 + 2H 2 = C 2 H 2 -f Ca(OH) 2 
Calcium Carbide Water Acetylene Calcium Hydroxide 

Calcium carbide is used to generate acetylene gas. This 
gas burns with a brilliant flame, and is coming into general 
use as an illuminant. Owing to its action with water, 

Light, Heat, Electricity, and Chemical Action. 117 

calcium carbide is packed and sold in air-tight cans (see 

Carborundum is a compound of silicon and carbon, hav- 
ing the composition SiC. It is made in the electric furnace 
by fusing sand (silicon dioxide, SiO 2 ), coke, saw.dust, and 
common salt. The essential chemical change is repre- 
sented thus : 

SiO 2 + 3C = SiC + 2 CO 

Silicon Dioxide Carbon Carborundum Carbon Monoxide 
Carborundum is silicon carbide (or carbon silicide). It 
is a crystallized solid, varying in color from white to emerald 
green and is sometimes iridescent. It is extremely hard, 
being harder than ruby and nearly as hard as diamond. 
Hence it is made into grinding wheels, whetstones, and 
polishing cloths. Over three million pounds were made at 
Niagara Falls in 1902, and the output is constantly 

Carborundum is a good conductor of heat. Its specific gravity is 
about three. Acids have no action upon it, but it is decomposed by 
fusing with potassium hydroxide and other alkalies. 

Carborundum is manufactured in a huge electric furnace, shown in 
Figure 18. It is an oblong box of bricks with permanent ends and loosely 
built sides. Each end is provided with a heavy metal plate. The wires 
for the electric current are attached to the outer ends of these plates, 
while the huge carbon electrodes fit into the inner ends, and project into 
the furnace. A cylinder of granulated coke makes an electrical connec- 
tion between the electrodes. In this furnace the rnixture is not heated 
by an electrical arc, but by the resistance of the carbon core to the pas- 
sage of the powerful current of electricity. The chemical change, as in 
the manufacture of calcium carbide, is due solely to heat. The current 
is passed through the mixture for about eight hours. When the opera- 
tion is over and the furnace is cool, the side walls are pulled down, and 
the carborundum is removed. The purest grade is found around the 
core. It is crushed, treated with sulphuric acid to remove the impurities, 
washed, dried, and graded according to the size of the particles. 


Descriptive Chemistry. 

Artificial Graphite is formed in the manufacture of 
carborundum. It is also made by heating a certain grade 
of anthracite coal in an electric furnace. It is extensively 
used in making electrodes for electric furnaces. Over 
800,000 Ib. were manufactured in 1902 at Niagara Falls. 
Graphite is a form of carbon (see Graphite). 

Light, Heat, Electricity, and Chemical Action. 119 


The Relation between Electricity and Chemical Action 

has always been a fascinating subject. Volta constructed 
his voltaic pile about 1800. This was one of the first, per- 
haps the first, source of an electric current. In May, 1800, 
Nicholson and Carlisle decomposed water into hydrogen 
and oxygen by an electric current obtained from a thermo- 
pile. In the same year Cruikshank obtained lead and 
copper from solutions of their salts. And in 1807 Davy 
isolated the elements, sodium and potassium, by passing an 
electric current (obtained from a large battery) through 
fused caustic soda and caustic potash respectively. From 
that time until the present day, the relation between elec- 
tricity and chemical action has engaged the attention of 
chemists. And their labors have built up a branch of 
chemistry called electrochemistry, which 
has recently attained considerable com- 
mercial importance. 

The Voltaic (or Galvanic) Cell in its simplest 
form consists of two metals connected by a wire 
and dipped into a liquid which will interact with 
one of the metals (Fig. 19). Copper, zinc, and 
water containing sulphuric acid may be used as 
an illustration. When the connected metals are FlG I9 ._voltaic cell, 
put into the acid, the zinc slowly disappears and 

hydrogen bubbles appear on the copper. Further examination would 
show that the zinc and sulphuric acid interacted, forming zinc sulphate. 
The chemical change is the one already described under hydrogen, and 
may be represented thus : 

Zn + H 2 S0 4 H 2 + ZnS0 4 

Zinc Sulphuric Acid Hydrogen Zinc Sulphate 

The connecting wire becomes electrified and exhibits the effects of an 
electric current, viz., it becomes warm, it makes a magnetic needle move, 

I2O Descriptive Chemistry. 

and a shower of sparks is produced if the wire is cut and one end is 
drawn down a file while the other is held firmly upon it. The source of 
the electric current is obviously the chemical action between the acid and 
zinc. The copper is necessary, otherwise the product of the chemical 
action would be merely heat. Carbon is often used in place of copper, 
and other liquids instead of sulphuric acid. The liquid chosen, how- 
ever, must be one that will interact with zinc or its substitute. Several 
cells joined together form an electric battery. For many years the 
battery was the chief source of the electric current. And it is now used, 
especially for ringing telephone, house, fire alarm, and signal bells, and 
in operating the telegraph. The dynamo is now widely used to generate 
powerful currents of electricity. 

Electrochemical Terms. Faraday (1791-1867) investi- 
gated electrochemistry about 1834, and introduced many 
terms in common use. He called the decomposing process 
electrolysis, and the decomposable liquid the electrolyte ; the 
wire by which the current entered he called the anode ; 
and that by which it escaped, the cathode. " Finally," he 
says, ^require a term to express those bodies which pass 
to the electrodes) I propose to distinguish such bodies by 
calling those anions which go to the anode of the decom- 
posing body ; and those passing to the cathode, cations ; 
and when I have occasion to speak of these together, I 
shall call them ions. Thus, chloride of lead is an electro- 
lyte, and when electrolyzed evolves the two ions, chlorine 
and lead, the former being an anion and the latter a 
cation." These terms are so used to-day, but they demand 
a broader definition. Electrolysis is the series of chemi- 
cal changes caused by the passage of an electric current 
through a dissolved or fused (i.e. melted) compound. The 
compound thus decomposed is an electrolyte. The metallic 
or carbon rods which conduct the current of electricity to 
and from the electrolyte are called the poles, or better, the 
electrodes. Electrodes are usually made of platinum, cop- 

Light, Heat, Electricity, and Chemical Action. 121 

per, zinc, mercury, or hardened carbon ; they may have 
any shape rod, wire, sheet, plate, box, crucible ; and they 
may also be solid, liquid, or powder, as well as fixed or 
movable. The electrodes are connected by wires with the 
source of the electric current, and serve as "doors" to 
quote Faraday again for the current to flow into and out 
of the electrolyte and through the wire connecting the 
electrodes. We speak of a " current" of electricity and 
of electricity as " flowing," although we do not know the 
nature of electricity, nor do we mean really that it flows, like 
a river, only in one direction. It is customary to speak of 
the current as entering the electrolyte by the anode or 
positive electrode and leaving by the negative electrode 
or cathode. The anode is the electrode that is often con- 
sumed or worn away, either mechanically or chemically. 
But solids are often deposited upon the cathode, as will 
soon be described, ^ons are those parts of the decomposed 
electrolyte which are believed to be material carriers of elec- 
tricij^y Aw comes from a Greek word which means wander- 
ing or migrating. And a cation is that ion which moves 
down or along with the current of electricity to the cathode 
where it is separated, deposited, or modified; while an anion 
is that ion which moves upward or against the current to the 
anode, where it likewise appears in various forms. Anions 
are electro-negative ions, but cations are electro-positive 
ions. Metallic ions are cations ; hence metals are deposited 
at the negative electrode or cathode. Non-metallic ions 
are usually anions, therefore oxygen, chlorine, and their 
oxides and hydroxides appear at the anode. Hydrogen is 
electro-positive. In general, metals are electro-positive, 
and non-metals (except hydrogen) are electro-negative. 
Ions follow the law of electric attraction and repulsion, 
viz., ions with the same kind of electrification repel each 


Descriptive Chemistry. 

FIG. 20. Electrolytic cell. A 
and C are the electrodes, R is the 
electrolyte, B or D is the battery 
or dynamo. 

other, and those with unlike kinds attract. Hence the 
electro-positive cations move toward the electro-negative 
cathode, and the electro-negative anions move toward the 
electro-positive anode. Ions are further described under 

lonization (see below). An elec- 
trolytic cell is the apparatus in 
which electrolysis takes place 
(Fig. 20). Its parts are analo- 
gous to the voltaic cell. There 
must be a containing Vessel, the 
two electrodes, and the electro- 
lyte. The vessel may have any 
desired shape, and is made of 
material which will resist the corrosive action of the electro- 
lyte or which will withstand a high temperature. Unlike the 
voltaic cell, the electrolytic cell generates no electric current ; 
it receives the current from a dynamo or a battery. Elec- 
trolysis is accomplished on a large scale in electrolytic cells. 

Illustrations of Electrolysis. Electrolysis may be 
simple, but it is usually very complex. Two illustrations 
will be given. When two platinum electrodes are put into 
melted zinc chloride and a current of electricity is passed, 
zinc is deposited at the cathode, and chlorine gas is liber- 
ated at the anode. This is a simple instance of electrolysis. 
But when an aqueous solution of sodium chloride is electro- 
lyzed, the action is different. Theoretically, the products 
should be sodium and chlorine, but they are hydrogen, 
sodium hydroxide, and chlorine. The sodium separated at 
the cathode immediately interacts with the water to form 
hydrogen and sodium hydroxide. Furthermore, unless the 
chlorine and sodium hydroxide are removed, they will 
interact to form compounds of chlorine, which vary in 
composition with the temperature, etc. 

Light, Heat, Electricity, and Chemical Action. 123 

The Electrolysis of Water is more complex than is ordinarily sup- 
posed. Strictly speaking, it is the sulphuric acid, and not the water, 
that is electrolyzed. Perfectly pure water does not conduct electricity, 
and is consequently not decomposed by it. But since the same amount 
of sulphuric acid is always present, no matter how long the action con- 
tinues, it is customary to speak of the total change as the electrolysis of 
water. The hydrogen and oxygen gases, which collect at the cathode 
and anode respectively, are merely the end products of a series of 
changes. Small quantities of ozone and hydrogen dioxide are also 

Faraday's Law. In his study of electrolysis, Faraday found that a 
measured quantity of electricity liberated different but definite amounts 
of the chemical elements. For example, the current which liberated 
i gm. of hydrogen also liberated 8 gm. of oxygen, 35.5 gm. of chlorine, 
108 gm. of silver, 31.7 gm. of copper, and so on. These numbers are 
identical with the chemical equivalents of these . elements (compare 
Equivalents, Chapter IX) . Faraday called them electrochemical equiv- 
alents, to emphasize their chemical and electrical relationship. But the 
term electrochemical equivalent now means, however, the weight of an 
element deposited or liberated by a current of a certain arbitrary value 
(i ampere in I second) . For example, the electrochemical equivalent of 
hydrogen is 0.000010441 gm., of oxygen is 0.00008287, and sometimes 
0.00016574, of copper is 0.0003294, and sometimes 0.0006588, of silver is 
0.001118. This general relation is often stated as Faraday's Law, 
thus : 

When the same quantity of electricity acts upon different electrolytes, 
the ratio between the quantities of liberated products is the same as 
between their chemical equivalents. 

Faraday also showed that the amount of decomposition the chem- 
ical work, we might say is proportional to the total amount of elec- 
tricity used. It makes no difference whether the current is strong or 
weak, nor whether the time of its flow is long or short. A certain 
quantity of electricity will do so much chemical work no more and 
no less. Thus a given quantity of electricity passed through copper 
sulphate solution always deposits the same weight of copper at the 
cathode. These two principles of Faraday are at the foundation of all 
electrochemical industries. Their importance can hardly be over- 

124 Descriptive Chemistry. 

Industrial Applications of Electrolysis. The earliest 
industrial application of electrolysis was in electrotyping 
and electroplating. These operations consist in depositing 
a thin film of metal upon a surface. They are fundamen- 
tally the same, though copper is the only metal used for 
producing electrotypes. Electrotypes are exact repro- 
ductions of the original objects. The process of electro- 
typing is substantially as follows : the page of type, or 
the woodcut, is first reproduced in wax or plaster. This 
exact impression is next covered with powdered graphite 
to make it conduct electricity. The coated mold is then 
suspended as the cathode in an acid solution of copper 
sulphate ; the anode is a plate or bar of copper. When 
the current is passed, electrolysis occurs ; copper is dis- 
solved from the anode and deposited upon the mold in a 
film of any desired thickness. The exact copper copy is 
stripped from the mold, backed with metal and mounted on 
a wooden block, and used instead of the type or woodcut 
itself. By this process exact copies of expensive wood 
engravings can be cheaply reproduced, and type can be 
saved from the wear and tear of printing. Most books, 
magazines, and newspapers are now printed from electro- 
types. The process of electroplating differs from elec- 
trotyping in only one essential, viz., in electroplating, the 
deposited film is not removed from the object. The object 
to be plated is carefully cleaned and made the cathode ; 
the anode is a bar or plate of the metal to be deposited. 
When the current passes through the system, the metal is 
firmly deposited upon the object. The electrolysis would 
take place, of course, if any anode were present ; but anodes 
of the metal to be deposited are usually used to prevent 
the solution or " bath " from weakening. They accom- 
plish the purpose by replenishing the solution with metal 

Light, Heat, Electricity, and Chemical Action. 125 

as fast as it is removed and deposited upon the cathode. 
Silver, nickel, and gold are the usual metals used in 
electroplating (see these metals). 

Electroplating and electrotyping have been done since 
about 1840. It is only within the last ten or fifteen years, 
however, that the electric current has been profitably 
applied in many industries. But during this time the 
development of electrochemistry has been very marked. 
The largest of these industries is the refining of copper. 
The process is similar to that described under electro- 
typing. Other metals, such as gold, silver, and lead, are 
extracted from their ores and purified by electricity, though 
the older processes are still used. All the aluminium, mag- 
nesium, and sodium of commerce are now manufactured by 
passing an electric current through their fused compounds. 
Nearly all the domestic potassium chlorate and much of 
the caustic soda are made by electricity. The same is 
true of barium compounds and many other chemicals. 
These electrochemical processes will be fully discussed in 
the appropriate places. 

The Theory of Electrolysis. Many theories have been 
proposed to explain electrolysis. According to the theory 
now generally held, electrolysis is not the splitting or tear- 
ing apart of molecules by the electric current. It is the 
carrying of electricity from one electrode to the other by 
ions. Dissolved or fused compounds are more or less dis- 
sociated into ions before the current of electricity is intro- 
duced, /and the current flows simply because the ions are 
there to carry it. Since these ions are charged with elec- 
tricity, the dissociation is called electrolytic dissociation 
or ionization. Ions are not atoms, but electrically charged 
atoms or groups of atoms. Thus, when sodium chloride is 
dissolved in water, much of the salt dissociates into the 

126 Descriptive Chemistry. 

ions, sodium and chlorine ; the sodium ions are charged 
positively, and the chlorine atoms negatively. Now, when 
an electric current is passed into the solution, the ions 
move toward their proper electrodes, carrying the electric 
charges with them. In brief, the current sorts the ions, 
which in turn migrate with their charges. When the ions 
reach their respective electrodes, they give up their electric 
charges and assume their normal conditions. Thus, the 
positive sodium ions give up their charges at the negative 
electrode, or cathode, and become sodium atoms. The 
latter interact with water to form hydrogen and sodium 
hydroxide. Similarly, the negative chlorine ions give up 
their charges at the positive electrode, or anode, and 
become neutral atoms, which at once unite to form chlo- 
rine molecules. 

Electrolysis and Solution. According to the above 
theory, the properties of many water solutions are closely 
related to the phenomena of electrolysis. For many years 
it was believed that a dissolved substance was distributed 
unchanged throughout the solvent. It was also believed 
that certain dissolved substances combined in part with 
the water a view held to-day. The first real step toward 
a settlement of the problem was taken when the electri- 
cal conductivity of solutions was compared. Experiments 
show that the electrical conductivity of solutions varies 
between wide limits. Water itself is practically a non- 
conductor, a sugar solution is a very poor conductor, while 
solutions of most acids, bases, and salts are excellent con- 
ductors. Water solutions, therefore, are of two kinds : 
(i) those which conduct electricity, and (2) those which 
do not, or only very slightly. But we have already seen 
that the first class consists of electrolytes. Hence, two 
things are believed about water solutions: (i) that when 

Light, Heat, Electricity, and Chemical Action. 127 

acids, bases, and salts are dissolved in water, they are dis- 
sociated into ions, and (2) that when sugar and similar 
substances are dissolved in water they dissociate very 
slightly or not at all. The amount of dissociation depends 
largely upon the relative amounts of solute and solvent, 
i.e. upon the dilution of the solution. The dissociation is 
slight in concentrated solutions, but increases as the dilu- 
tion increases. Not all acids, bases, and salts dissociate to 
the same degree. The percentage of dissociation of some 
of these compounds in solutions of a certain strength and at 
the same temperature (i 8 C.) is given in the following 




Hydrochloric acid 



Potassium chloride 


Potassium nitrate 

/ j 

Potassium hydroxide 


Sodium hydroxide 


Numerous facts support the theory of ionization. (i) 
Varying electrical conductivity has already been mentioned. 
(2) It has long been known that solutions boil at a higher 
temperature and freeze at a lower temperature than pure 
water. A fresh-water river, for example, freezes before 
the ocean, and water containing considerable mineral mat- 
ter boils at a higher temperature than pure drinking water. 
It is generally true that a dissolved substance raises the 
boiling point and loivers the freezing point of a given solu- 
tion. Now, when weights of substances proportional to 

128 Descriptive Chemistry. 

their molecular weights are dissolved in the same volume of 
water, the boiling point of each solution is raised the same 
number of degrees and the freezing point is lowered the 
same number of degrees. These facts are now applied 
experimentally to determine molecular weights. In many 
cases the molecular weights thus found agree with the 
values obtained by other methods. Thus, if X is the 
depression produced by a one per cent solution of sugar, 
and Y the depression produced by a one per cent solution 
of urea, the following proportion may be written, because 
the depressions of the freezing points are inversely propor- 
tional to the molecular weights 

Y: X : : mol. wt. of sugar : mol. wt. of urea. 

The molecular weight of sugar is known to be 342, and 
from the proportion the molecular weight of urea is 60, 
which agrees with that found by other methods. This 
method is applicable to many compounds and is helpful 
in deciding whether a molecular weight is a given number 
or its multiple. There is a marked disagreement to this 
rule, however, in the case of solutions of acids, bases, and 
salts. That is, electrolytes are exceptions. In some 
instances the molecular weight is only half that found 
by other methods. Thus, the molecular weight of sodium 
chloride was found to be about 30, instead of 58.5 the 
correct molecular weight. Hence, it is believed that the 
solutions of acids, bases, and salts contain ions which act 
like molecules in their effect upon the freezing and boiling 
points of solutions. The behavior of acids, bases, and 
salts in solution led the Swedish chemist Arrhenius, in 1887, 
to extend the ideas of Faraday and to propose the present 
theory of solution. 

Light, Heat, Electricity, and Chemical Action. 129 

Application of the Theory of lonization. Many ob- 
scure facts of chemistry become intelligible when inter- 
preted by the theory of ionization. (i) Ordinary tests are 
tests for ions. For example, all chlorides in solution have 
the same test. That is, they all interact with silver nitrate 
in solution, because all have chlorine ions in the solution. 
Similarly, all soluble sulphates interact with barium chlo- 
ride in solution, because all sulphates have SO 4 ions in the 
solution. Both silver chloride and barium sulphate are 
insoluble, and are removed from the solution as precipi- 
tates. A complete illustration will make this fact clearer. 
The silver nitrate and sodium chloride solutions before 
mixing consist largely of the ions of silver, NO 3 -group, 
sodium, and chlorine. When mixed, the ions of silver 
and chlorine unite to form silver chloride, which is in- 
soluble and hence not ionized ; the solution still contains 
ions of sodium and of the NO 3 -group. On the other 
hand, if solutions of potassium chlorate and silver nitrate 
are mixed, no silver chloride is formed, because no chlo- 
rine ions are available. Potassium chlorate dissociates 
into ions of potassium and C1O 3 . Equations are often 
used to express ionization. Thus, the ionic equation for 
the interaction of sodium chloride and silver nitrate is 

Na + Cl 4 Ag + NO 8 = AgCl 4- Na 4 NO 3 . 

(2) lonization explains the General Properties of Acids, 

Bases, and Salts. Acids in solution turn litmus red, be- 


cause their solutions contain hydrogen ions (H). Simi- 
larly, bases turn litmus blue, because their solutions contain 

hydroxyl ions (OH). But solutions of neutral salts con- 
tain neither hydrogen nor hydroxyl ions, hence they do 
not affect litmus. The above principles can be readily 

i jo Descriptive Chemistry. 

extended to cover acid and basic salts. The other general 
properties of acids and bases are believed to be due to the 
above causes. (3) Neutralization, interpreted by the ionic 
theory, is fundamentally the union of hydrogen and hy- 
droxyl ions to form molecules of water. Suppose hydro- 
chloric acid and potassium hydroxide are mixed. The 
solution at first contains the hydrogen, chlorine, potassium, 
and hydroxyl all as ions. But the hydrogen and hy- 
droxyl immediately unite to form water, leaving the po- 
tassium and chlorine ions in the solution. This solution 
is thus rendered neutral by the removal of the hydrogen 
ion its acid constituent and of the hydroxyl ion its 
basic constituent. The ionic equation expressing the 
neutralization of potassium hydroxide by hydrochloric 
acid is 

K + OH + H + Cl = K + Cl + H 2 O. 
The potassium and chlorine ions remain free and un- 
combined until the solution is evaporated. As the con- 
centration increases, the ions unite until nothing remains 
except* the neutral salt potassium chloride. 

Neutralization, therefore, as interpreted by the ionic theory, is essen- 
tially a union of hydroxyl and hydrogen ions. This view is supported 
by much experimental evidence. For example, the heat of neutraliza- 
tion produced by the interaction of equivalent quantities of strong 
acids and bases is approximately the same. 


1. What transformations of energy accompany chemical action? 
Illustrate your answer. 

2. State and illustrate the law of the conservation of energy. 

3. Discuss the relation of light to chemical action. Give popular 
and scientific illustrations of (a} the production of chemical action by 
light, and (#) production of light by chemical action. 

4. Define and illustrate (a} calorie, (6) thermal equation, (c} heat 
of formation, (d} exothermic, (e) heat of decomposition, (/) endothermic. 

Light, Heat, Electricity, and Chemical Action. 131 

5. Give several illustrations of the production of (a} heat by chemi- 
cal action, and (b) vice versa. 

6. When an electric spark is passed through a mixture of two vol- 
umes of hydrogen and one volume of oxygen, what is the result? Is it 
due directly to electricity or to heat? 

7. Define and illustrate kindling temperature. 

8. Name several sources of heat. How may electricity be used as 
a source of heat ? 

9. Describe Moissan's electric furnace. Why is it so efficient? Is 
its effect thermal or electrical ? State some results produced by Moissan 
with this furnace. Has the electric furnace any industrial use ? Where ? 

10. What is calcium carbide? How is it made? State the equation 
for the reaction. What are its properties ? For what is it used? 

11. What is carborundum? How is it made? State the equation 
for the reaction. What are its properties and uses? 

12. What is artificial graphite? How is it made? For what is it used? 

13. Give several illustrations of the production of (a) electricity by 
chemical action, and (b) vice versa. 

14. State briefly the first chemical changes which were produced by 

15. Describe a simple voltaic cell. Why is it so called? What is 
the source of the electric current manifested by the cell? What is an 
electric battery ? For what is it used ? 

1 6. Define and illustrate (a) electrolysis, (b) electrolyte, (c} elec- 
trode, (W) anode, (e) cathode, (/) ions, (g) anion, (h} cation, (*) posi- 
tive electrode, (/) negative electrode, () ionization. 

17. Where are (a) anions and (b) cations liberated? 

1 8. Describe an electrolytic cell. How does it differ from a voltaic 
cell ? For what is it used ? 

19. Describe the electrolysis of (a) zinc chloride, (b} sodium chloride, 
(c) water. 

20. State and illustrate Faraday's law. 

2 1 . Give a brief account of Faraday's contribution to electrochemistry. 

22. Describe the process of (a) electrotyping and (b} electroplating. 

23. State some industrial applications of the electric current. 

24. What is the theory of electrolysis? What is the present theory 
of solution in water? What is the theory called? Why? What facts 
support it? 

25. Define and illustrate an ionic equation. 

Descriptive Chemistry. 


1. Calculate the percentage composition of (a) water, () magnetic 
oxide of iron (Fe 3 O 4 ), (c} crystallized sodium carbonate (Na.,CO 3 . 
ioH 2 O). 

2. If a certain current of electricity deposited 31.7 gm. of copper, 
how much (a) silver, (&) aluminium, and (c) magnesium would it 
deposit ? 

3. If a certain current of electricity deposited 2 kg. of copper, how 
much silver would it deposit? 

4. How much calcium carbide can be made (theoretically) from a 
ton of lime? (Equation is 3 C + CaO = CaC 2 + CO or 36 + 56 = 
64 + 28.) 

5. How much carborundum can be made (theoretically) from a ton 
of sand (SiO 2 ) ? (Equation is SiO 2 + 3 C = SiC + 2 CO or 60 + 36 
= 40 + 56.) 

6. Calculate the percentage composition of (#) carborundum and 
(&) calcium carbide. 


CHLORINE is an important element, and its compounds 
are useful, especially hydrochloric acid, sodium chloride, 
and bleaching powder. 

Occurrence. Free chlorine is never found in nature, 
because it combines so readily with other elements. But 
in combination it is widely distributed, since it is one of 
the components of common salt, or sodium chloride. 
Many compounds of chlorine with potassium, magnesium, 
and calcium are found in the deposits at Stassfurt in Ger- 
many (see these metals). The salts found in sea water 
contain about 2 per cent, and the earth's crust contains 
about o.oi per cent of chlorine. Silver chloride "horn" 
silver is mined as an ore in the United States and 

Preparation. Chlorine is prepared in the laboratory 
by heating a mixture of manganese dioxide and hydro- 
chloric acid. This method was used by Scheele, who 
discovered the gas in 1774. The equation for the prepa- 
ration of chlorine is 

MnO 2 + 4HC1 = C1 2 + MnCl 2 + 2 H 2 O 
Manganese Hydrochloric Chlorine Manganese Water 

Dioxide Acid Bichloride 


This is an oxidizing process, since the hydrogen of the 
hydrochloric acid is oxidized to water, although only part 
of the chlorine of the acid is obtained free. 

134 Descriptive Chemistry. 

Sometimes chlorine is prepared in the laboratory by heating a mixture 
of manganese dioxide, sodium chloride, and sulphuric acid. This method 
is substantially the same as the other, since a mixture of sulphuric acid 
and sodium chloride yields hydrochloric acid. The simplest equation 
for this method of preparing chlorine is 

2 H 2 SO 4 + 2 NaCl + MnO 2 = C1 2 + Na,SO 4 + MnSO 4 -f 2 H 2 O 
Sulphuric Sodium Manganese Chlorine Sodium Manganese Water 
Acid Chloride Dioxide Sulphate Sulphate 

Other oxidizing substances besides manganese dioxide may be used, 
such as potassium chlorate (KC1O 3 ), potassium dichromate (K 2 Cr 2 O 7 ), 
and red lead (Pb 3 O 4 ). 

Chlorine is manufactured by several processes, all of 
which involve the same principle as the laboratory method. 

In the Deacon process, hydrochloric acid is oxidized by oxygen ob- 
tained from the atmosphere. A mixture of hydrochloric acid gas and 
air is heated to 500 C. and passed through iron tubes containing balls 
of clay or pieces of brick previously saturated with copper chloride. A 
series of complex reactions occurs which are not well understood. It is 
supposed that the copper chloride facilitates the formation of chlorine 
by continuously giving and taking this gas. The essential chemical 
change, however, is the oxidation of the hydrochloric acid, and it may 
be represented by the equation 

2HC1 -I- O C1 2 -f H 2 O 

Hydrochloric Acid Oxygen Chlorine Water 

In the Weldon process, an impure native manganese dioxide, known as 
pyrolusite, is treated with hydrochloric acid in large earthenware retorts 
or stone tanks heated by hot water or steam. When no more chlorine 
is liberated, the residue is mainly manganese dichloride. This " still- 
liquor" was formerly thrown away, but by the Weldon process it is 
changed into manganese compounds, which are used to prepare more 
chlorine (see Manganese Dioxide). 

Chlorine is also prepared on a large scale by the electrolytic process. 
Sodium chloride is decomposed by electricity in properly constructed 
cells, and the chlorine which is liberated at the anode is conducted off 
through pipes to the bleaching powder factory. Sodium hydroxide is 
produced at the same time, and the process will be described under this 

Chlorine and Hydrochloric Acid. 135 

Properties. Chlorine is a greenish yellow gas. Its 
color suggested the name chlorine (from the Greek word 
chloros, meaning greenish yellow), which was given to it 
by Davy about 1810. It has a disagreeable, suffocating 
odor, which is very penetrating. If breathed, it irritates 
the sensitive lining of the nose and throat, and a large 
quantity would doubtless cause death. It is heavier than 
the other elementary gases, and is about 2.5 times heavier 
than air. Hence it is easily collected by downward dis- 
placement, i.e. by allowing it to fall to the bottom of a 
bottle and thus fill the latter by displacing the air. 

A liter of dry chlorine at o C. and 760 mm. weighs 3.18 gm. 

Water dissolves chlorine. The solution is yellowish, 
smells strongly of chlorine, and is frequently used in the 
laboratory as a substitute for the gas. Chlorine water, as 
the solution is called, is unstable even under ordinary con- 
ditions, and must be kept in the dark. If the solution is 
placed in the sunlight, oxygen is soon liberated and hydro- 
chloric acid is formed. Intermediate changes doubtless 
occur ; but the simplest equation for the essential change 

is- Hp + C1 2 = 2HC1 + O 

Water Chlorine Hydrochloric Acid Oxygen 

Chlorine is much less soluble in a solution of sodium chloride, over 
which it is sometimes collected. It attacks mercury and cannot be col- 
lected over this liquid. 

Chlorine does not burn in the air, but many substances 
burn in chlorine. The metals antimony and arsenic, when 
sprinkled into chlorine, suddenly burst into flame, while 
phosphorus melts at first and finally burns with a feeble 
flame. If sodium, iron powder, brass wire, or other metals 
are heated and then put into chlorine, they burn ; the' 
sodium and iron produce a dazzling light and the brass 

136 Descriptive Chemistry. 

glows and emits dense fumes of whitish smoke. Chlorine 
combines readily with hydrogen. Hence, a jet of burning 
hydrogen when lowered into chlorine continues to burn, 
forming hydrochloric acid gas, which appears as a white 
cloud. The simplest equation for this change is 

H + Cl HC1 

Hydrogen Chlorine Hydrochloric Acid 

The attraction between chlorine and hydrogen is so great 
that many compounds of hydrogen are decomposed by 
chlorine. Thus, compounds containing hydrogen and 
carbon, such as illuminating gas, paraffin wax, and wood, 
burn in chlorine with a smoky flame. Chlorine does not 
combine directly with carbon, hence the flame consists 
largely of very fine particles of solid carbon. Similarly, 
a piece of glowing charcoal is extinguished by chlorine. If 
filter paper is saturated with warm turpentine (a compound 
of hydrogen and carbon) and put into a bottle of chlorine, 
a flame accompanied by a dense cloud of black smoke 
bursts from the bottle ; the chlorine withdraws the hydro- 
gen to form hydrochloric acid, while the carbon is left free. 
The power to bleach is the most striking and useful 
property of chlorine. This property depends upon the 
fact, already mentioned, that chlorine withdraws hydrogen 
and liberates free oxygen ; the latter then decomposes the 
coloring matter in the cloth or other material. Dry 
chlorine does not bleach. If an envelope on which the 
postmark, or a lead pencil mark, is still visible is placed 
in moist chlorine, these marks will not be bleached be- 
cause they are largely carbon ; but the writing ink, which 
is mainly a compound of hydrogen, carbon, and iron, will 
disappear. Litmus paper and calico are both bleached by 
moist chlorine. 

Chlorine and Hydrochloric Acid. 137 

Bleaching Powder is the source of the chlorine used in 
the bleaching industries. It is sometimes called "bleach," 
or " chloride of lime." It is a yellowish white substance 
having a peculiar odor, which resembles that of chlorine. 
When dry, it is a powder, but on exposure to the air, it 
absorbs water and carbon dioxide, becomes lumpy and 
pasty, and loses some of its chlorine. Acids like sulphuric 
and hydrochloric acid liberate from bleaching powder its 
" available chlorine," which varies from 30 to 38 per cent 
in good qualities. The equations for the interaction of 
acids and bleaching powder are usually written thus 

CaOCl 2 + H 2 SO 4 = Cl a + CaSO 4 + H 2 O 
Bleaching Powder Sulphuric Acid Calcium Sulphate 

CaOCl 2 + 2 HC1 = C1 2 -f CaCl 2 + H 2 O 
Hydrochloric Acid Calcium Chloride 

The composition of bleaching powder has been much discussed. 
The most reliable authority gives it the formula CaOCL,. When dis- 
solved in water, bleaching powder forms calcium hypochlorite (CaO 2 Cl 2 ) 
and calcium chloride (CaCl.,). 

Bleaching Powder is manufactured by the action of chlorine gas on 
lime. Lime (calcium oxide, CaO) is carefully slaked with water to 
form calcium hydroxide (Ca(OH). 2 ). This powder is sifted into a 
large absorption chamber made of iron, lead, or tarred brick until the 
floor is covered with a layer three or four inches deep. The chlorine 
enters at the top and settles slowly to the floor, where it is absorbed 
by the lime. 

The simplest equation for the formation of bleaching powder might 
be written 

Ca(OH) 2 + C1 2 CaOCl 2 + H 2 O 

Calcium Hydroxide Chlorine Bleaching Powder Water 

Bleaching. Immense quantities of bleaching powder 
are used to whiten cotton and linen goods and paper pulp. 
The pieces of cotton cloth as they come from the mill are 

138 Descriptive Chemistry. 

sewed end to end in strips, which are stamped at the 
extreme ends with some indelible mark to distinguish each 
owner's cloth. These strips, which are often several miles 
long, are drawn by machinery into and out of numerous 
vats of liquors and water, between rollers, and through 
machines, until they are snow-white and ready to be 
finished (i.e. starched and ironed) or dyed. The whole 
operation requires three or four days. 

The preliminary treatment consists in singeing off the downy pile 
and loose threads by drawing the cloth over hot copper plates or 
through a series of gas flames. The object of the remaining operations 
is threefold, (i) to wash out mechanical impurities, the fatty and resin- 
ous matter, and the excess of the different chemicals, (2) to remove 
matter insoluble in water, and (3) to oxidize the coloring matter by 
chlorine. The details of the process differ with the texture of the 
cloth and with its ultimate use. The threefold object above mentioned 
involves successively "liming," "souring," "chemicking," and "souring, 11 
interspersed with frequent washing. The "liming 11 consists in boiling 
the cloth in a large kier or vat with lime, the "souring 11 in wetting it 
with weak sulphuric or hydrochloric acid, and the " chemicking ?1 in im- 
pregnating it with a weak solution of bleaching powder. Often the cloth 
is boiled at a certain stage with resin and sodium carbonate. The 
^liming 11 removes the resinous and the fatty matter, the first "souring 11 
neutralizes traces of lime, and the second, which follows the "chem- 
icking, 11 liberates the chlorine in the fiber of the cloth. Frequent washing 
is absolutely necessary to remove the impure products of the chemical 
changes as well as the excess of lime and other alkali, acid, and chlo- 
rine. Should these be left, the cloth would be unevenly bleached and 
its fiber would be weak. The cloth is finally treated with an antichlor, 
such as sodium hyposulphite, which removes the last traces of chlorine. 

Bleaching is chemically an oxidizing process. The 
oxygen when it is liberated from water by chlorine is said 
to be in the nascent state. This means that the gas is 
exceedingly active, because it is not only uncombined, but 
just ready to unite with those elements for which it has 
great affinity. Hence this nascent oxygen literally tears 

Chlorine and Hydrochloric Acid. 139 

down complex colored substances and changes them into 
colorless compounds. The nascent state is aptly illustrated 
by bleaching because both the chlorine and the oxygen 
are in this active chemical condition. 

Chlorine Hydrate is formed by cooling chlorine water or by passing 
chlorine into ice water. It is a yellowish, crystalline solid, and in the 
air it decomposes quickly into chlorine and water. Its composition 
corresponds to the formula C1 2 10 H 2 O. 

Liquid Chlorine was first prepared by Faraday in 1823. A little 
chlorine hydrate was inclosed in one arm of a bent tube (Fig. 21), 
which was then sealed. By gently heating the tube, the chlorine hy- 
drate was decomposed into chlorine and water, 
but the chlorine, being unable to escape, was 
condensed to a liquid by the pressure inside the 
tube. The liquefaction is more easily accom- 
plished if one end is kept cold during the 
experiment. FlG - 21.- Bent tube for 

. . ,. ,, . .the liquefaction of chlo- 

At the ordinary pressure, chlorine gas be- r j ne 

comes liquefied, if its temperature is 34 C, 

while at a pressure of six atmospheres the temperature need be only 
o C. Liquid chlorine has a bright yellow color. It is a commercial 
article, and is stored and shipped in steel cylinders lined with lead. 
It is used in the laboratory to prepare chlorides, and industrially to 
extract gold. Solid chlorine has been obtained as a yellow crystalline 
mass by cooling the liquid to 102 C. 

Uses of Chlorine. Chlorine is used directly to prepare 
some of its compounds, the most important being bleaching 
powder. The latter is often used as a deodorizer and dis- 
infectant, since the liberated chlorine destroys putrefying 
matter by acting on it as on coloring matter. A solution 
of potassium hypochlorite (Javelle's water) or sodium hy- 
pochlorite (Labarraque's solution) is often used to remove 
fruit stains from cotton and linen goods. 

Chlorides are formed when chlorine combines with other 
elements, and they are in general stable compounds. 

140 Descriptive Chemistry. 

The simplest equations illustrating the combination of chlorine with 
metals and other elements are 

Na + Cl = NaCl 

Sodium Chlorine Sodium Chloride 

Sb + 3d = SbCl 3 

Antimony Antimony Trichloride 

Cu + C1 2 = CuCl 2 

Copper Copper Chloride 

P + 3 C1 = PC1 3 

Phosphorus Phosphorus Trichloride 

H + Cl = HC1 

Hydrogen Hydrochloric Acid 

Chlorides form an important class of compounds and they will be 
considered under the elements with which the chlorine combines. 
(See also Chlorides below.) 


Hydrochloric Acid is the most useful compound of 
chlorine. It is a gas, very soluble in water^ This solution 
has long been known as muriatic acid (from the Latin 
word muria, meaning brine). The term hydrochloric acid 
includes both the gas and its solution, but the solution is 
usually meant. 

The early chemists called the gas " spirit of salt." Priestley, who 
first prepared, collected, and studied the gas, called it " marine acid air." 
Both expressions emphasize its relation to salt (sodium chloride). 

Occurrence. The gas occurs free in volcanic gases. 
The solution is one constituent of the gastric juice of the 
stomach. Chlorides, which are salts of hydrochloric acid, 
are abundant in the earth's crust. 

Preparation. The gas is prepared in the laboratory 
by the method devised by Glauber in the seventeenth cen- 

Chlorine and Hydrochloric Acid. 141 

tury, viz., by heating sulphuric acid and sodium chloride. 
If the mixture is gently heated, the chemical change is 
represented thus 

Nad + H 2 S0 4 = HC1 + HNaSO 4 
Sodium Sulphuric Hydrochloric Acid Sodium 
Chloride Acid Acid Sulphate 

But at a high temperature the equation for the reaction 
2 NaCl + H 2 SO 4 - 2 HC1 + Na 2 SO 4 

In either case the gas is readily produced. It may be 
collected over mercury or, more easily, by downward dis- 
placement. The solution is prepared by passing the gas 
into water. 

That sodium sulphate is the other product of the chemical change at 
a high temperature may be shown by testing the heated residue as 
follows : (a) Dissolve a portion in water and add a few drops of barium 
chloride solution ; the immediate formation of the white, insoluble 
barium sulphate shows that the residue from the experiment must be 
a sulphate, (b} Burn a little of the residue on a platinum wire or 
piece of porcelain held in the Bunsen flame ; the intense yellow color 
immediately imparted to the flame shows that the residue contains 
sodium, (c) Hence the compound must be sodium sulphate. 

Commercial Hydrochloric Acid is manufactured in enor- 
mous quantities by the method used in the laboratory. 
A mixture of salt and sulphuric acid is moderately heated 
in a large hemispherical cast-iron pan, and the gas passes 
through an earthenware pipe into an absorbing tower ; the 
fused mass of acid sodium sulphate and salt is then sub- 
jected to a higher temperature, and the liberated gas passes 
by another pipe into the absorbing tower. These towers 
are tall and filled with coke or pieces of brick over which 
water trickles ; as the hydrochloric acid gas passes up the 
tower, it is absorbed by the descending water, and flows 

142 Descriptive Chemistry. 

out at the bottom of the tower as concentrated acid. The 
gas is usually cooled before it enters the towers. Some- 
times the gas passes through huge earthenware jars be- 
fore entering the towers. In these jars the gas and water 
are caused to flow constantly in opposite directions, thus 
insuring complete absorption. 

Hydrochloric acid gas is a by-product in the manufacture of sodium 
carbonate by the Leblanc process. The gas was formerly allowed to 
escape into the atmosphere, but since it destroyed vegetation and be- 
came a nuisance in other ways, a law was passed forbidding the manu- 
facturers to let it escape. Hence it became necessary to absorb the 
gas in water. The hydrochloric acid, which was once regarded as a 
waste product, is now the main source of profit, since competition has 
reduced the price of sodium carbonate (see Sodium Carbonate). 

Properties. Hydrochloric acid gas is colorless and 
transparent. When it escapes into moist air, it forms 
fumes which are really minute drops of a solution of the 
gas in the moisture of the air. It has a choking, sharp, 
pungent odor. The gas does not burn nor support com- 
bustion. It is about 1.25 times heavier than air, and may 
therefore be collected by downward displacement. 

One liter at oC. and 760 mm. weighs 1.61 gm. The gas can be 
liquefied at ioC. and 40 atmospheres pressure; while at i6C, the 
pressure need be only 20 atmospheres. 

The extreme solubility of hydrochloric acid gas in water is 
one of its most striking properties. One liter of water will 
dissolve about 500 1. of gas, if both are at o C. and 760 mm. 
At the ordinary temperature about 450 1. of gas dissolve in 
i 1. of water, and as the temperature rises the solubility 
decreases. The solution is the familiar hydrochloric acid. 
The gas readily escapes, hence the acid forms fumes when 
exposed to air. Pure hydrochloric acid is a colorless liquid. 
The commercial acid has a yellow color, usually due to iron 

Chlorine and Hydrochloric Acid. 143 

compounds, but sometimes to organic matter or to dissolved 
chlorine. It also contains other impurities. Like most 
acids, it reddens blue litmus, and gives up its hydrogen 
when added to metals. 

The strongest acid contains about 42 per cent (by weight) of the 
gas, and its specific gravity is i .2. When the strong acid is heated, the 
gas is evolved until the solution contains about 20 per cent of the acid, 
and then the liquid boils at i ioC. without further change. The dilute 
acid, on the other hand, loses water until the same conditions prevail. 

Composition of Hydrochloric Acid Gas. In 1810, Davy showed 
that hydrochloric acid gas (which had been regarded as an oxygen 
compound) contained only chlorine and hydrogen. Many facts lead 
us to conclude that hydrochloric acid gas is composed of hydrogen and 
chlorine in such a ratio that its composition is represented by the for- 
mula HC1. (i) Hydrogen burns in chlorine, and the only product is 
hydrochloric acid gas. (2) When hydrochloric acid is decomposed 
by an electric current, equal volumes of hydrogen and chlorine are 
evolved. (3) When a mixture of equal volumes of hydrogen and 
chlorine is exposed to the direct sunlight or to the action of an electric 
spark, the gases combine with an explosion, and hydrochloric acid gas 
is formed with no residue. Furthermore, the volume of the resulting 
gas equals the sum of the volumes of hydrogen and chlorine used. 
(4) When a given volume of dry hydrochloric acid gas is treated with 
sodium amalgam, the chlorine is withdrawn by the sodium in the amal- 
gam, and a volume of hydrogen remains which is half the original vol- 
ume. (5) No derivative of hydrochloric acid is known which contains 
less hydrogen, or less chlorine in a molecule. (6) The ratio by weight 
in which hydrogen and chlorine combine is 1:35.45. Hence, the 
lowest molecular weight of hydrochloric acid is 36.45, a number which 
has been verified by several different methods. 

Uses of Hydrochloric Acid. Vast quantities are used 
to prepare chlorine for the manufacture of bleaching pow- 
der. Various chlorides are prepared from it, and it is one 
of the common acids used in chemical laboratories. 

Chlorides are formed by the direct addition of chlorine 
to metals, as we have seen. They are also formed when 

144 Descriptive Chemistry. 

metals, their oxides, or hydroxides are added to hydro- 
chloric acid. The following equations illustrate this gen- 
eral fact : 

Zn + 2 HC1 = ZnCl 2 + H 2 

Zinc Zinc Chloride 

ZnO + 2 HC1 = ZnCl 2 + H 2 O 

Zinc Oxide Zinc Chloride 

Zn(OH) 2 + 2 HC1 - ZnCl 2 + 2 H 2 O 

Zinc Hydroxide Zinc Chloride 

They are also formed by adding other salts to hydro- 
chloric acid. 

Molecules of chlorides may contain several atoms of chlorine. 
Occasionally the- name of the compound indicates this fact, e.g. manga- 
nese dichloride (MnQ 2 ), antimony trichloride (SbCl ;5 ), phosphorus 
trichloride and pentachloride (PC1 3 and PCI-)- If a metal forms two 
chlorides, the two are distinguished .by modifying the name of the 
metal. The one containing the smaller proportion of chlorine ends in 
-ous, the one containing the larger ends in -ic. Thus, mercurous chlo- 
ride is HgCl, but HgCl 2 is mercuric chloride. Similarly, we have fer- 
rous chloride, FeCl 2 , and ferric chloride, FeCl 3 . 

The Test for Hydrochloric Acid and Chlorides. Most 

chlorides are soluble in water. Those of lead, silver, and 
mercury (-ous) are not. If silver nitrate is added to hydro- 
chloric acid, or to the solution of a chloride, a white, curdy 
precipitate of silver chloride is formed, which (a) is insol- 
uble in nitric acid, but soluble in warm ammonium hydrox- 
ide, and () turns purple in the sunlight. The invariable 
formation of silver chloride is the test for hydrochloric 
acid and soluble chlorides. Hydrochloric acid gas also 
forms dense white clouds of ammonium chloride in the 
presence of ammonia gas. 

Chlorine and Hydrochloric Acid. 145 

Miscellaneous. The acids of chlorine are tabulated under ACIDS. 
The compounds of chlorine with sodium, potassium, magnesium, and 
calcium are described under these metals. 

Aqua regia, of which chlorine is one constituent, is discussed in 
Chapter XII. 


1. What is the symbol of chlorine ? What useful compounds con- 
tain this element ? 

2. How is chlorine prepared in the laboratory ? Give one equation 
for its preparation. Describe Deacon^ process for manufacturing 

3. Who discovered chlorine ? Who named it, when, and why ? 

4. Summarize the physical properties of chlorine. How can it be 
quickly distinguished from the gases previously studied ? 

5. Summarize the chemical properties of chlorine. Compare it 
with oxygen. Describe fully its action with hydrogen. 

6. Define (a) downward displacement, (b} available chlorine, 
(V) antichlor. 

7. Develop the topics : (a) nascent state, (<$) chlorine water, (V) chlo- 
rine hydrate, (ti ) liquid chlorine, (e) chlorine is an oxidizing agent. 

8. What is bleaching powder ? How is it made ? What are its 
chief properties ? Describe the operation of bleaching. What is the 
chemistry of bleaching ? 

9. What is (a) "bleach," () muriatic acid, (V) chloride of lime, 
(d) "salt, 11 (e) "lime," (/) commercial hydrochloric acid ? 

10. What are chlorides ? Name five. How can they be formed ? 
Give the formula of sodium chloride. Why cannot chlorine be collected 
over mercury ? 

11. What is hydrochloric acid ? How is it prepared in the labora- 
tory ? Give the equations for its preparation. How is it prepared 
industrially ? 

12. Summarize the chief properties of hydrochloric acid gas. Of the 
acid, as the term is usually used. What happens when hydrochloric 
acid is boiled ? 

13. What is the evidence that the formula of hydrochloric acid gas 
is HC1 ? 

14. For what is hydrochloric acid used ? State the test for hydro- 
chloric acid and soluble chlorides. 

146 Descriptive Chemistry. 

15. Give a brief account of Faraday's work on chlorine. Of Davy's 

1 6. Why is chlorine never found free ? 


1. One equation for the preparation of chlorine is 

4HC1 + MnO 2 = C1 2 + MnCI 2 + 2H 2 O 
146 + 87 =71 + 126 + 36 

(0) How many grams of chlorine can be made from 247 gm. of man- 
ganese dioxide ? () Name all the products. 

2. How much sodium chloride is needed to prepare a kilogram of 
hydrochloric acid gas ? 

3. How many grams of manganese dioxide are necessary to ^prepare 
100 gm. of chlorine from hydrochloric acid. 

4. A bottle of chlorine water was exposed to the sunlight until 
all the chlorine disappeared, (a) What two products were formed ? 
(^) Write the equation for the reaction. (c} What weight of chlorine 
gas is necessary to form 20 gm. of the gaseous product ? (d) What 
volume of chlorine is necessary to form 20 gm. of the other product ? 

5. Calculate the percentage composition of (a) hydrochloric acid 
gas, (b) sodium chloride, (c) silver chloride (AgCl), (d) potassium 
chloride (KC1). 


THE most important compounds of nitrogen are am- 
monia (NH 3 ), nitric acid (HNO 3 ), and compounds related 
to them. Many animal and vegetable substances essential 
to life are compounds of nitrogen. 


The term ammonia includes both the gas and its solu- 
tion in water, though the latter is more accurately called 
ammonium hydroxide. 

Formation of Ammonia. When vegetable and animal 
matter containing nitrogen decays, the nitrogen and hydro- 
gen are liberated in combination, as ammonia. The odor 
of ammonia is often noticed near stables. If animal sub- 
stances containing nitrogen are heated, ammonia is given 
off. The old custom of preparing ammonia by heating 
horns and hoofs in a closed vessel, i.e. by dry distillation, 
gave rise to the term "spirits of hartshorn." Soft coal 
contains compounds of nitrogen and of hydrogen, and when 
the coal is heated to make illuminating gas, one of the prod- 
ucts is ammonia. 

Preparation. Ammonia gas is prepared in the labora- 
tory by heating ammonium chloride with an alkali, usually 
slaked lime. The reaction may be represented thus 


148 Descriptive Chemistry. 

2NH 4 C1 + Ca(OH) 2 = 2NH 3 + CaCl 2 + 2 H 2 O 
Ammonium Slaked Ammonia Calcium 

Chloride Lime Gas Chloride 

107 + 74 =34 + 111 + 36 
The gas is usually collected by upward displacement, i.e. 
by allowing the gas to flow upward into a bottle and dis- 
place the air. The solution is prepared by conducting the 
gas into water. 

The main source of the ammonia of commerce is the ammoniacal 
liquor or gas liquor of the gas works. The gases which come from the 
retorts in which the coal is heated are passed into water, which absorbs 
the ammonia and some other gases. This impure gas liquor is treated 
with lime to liberate the ammonia, which is absorbed in tanks contain- 
ing hydrochloric acid or sulphuric acid. This solution upon the addi- 
tion of an alkali gives up its ammonia, which is dissolved in distilled 
water, forming thereby the ammonium hydroxide or aqua ammonia 
of commerce. 

Ammonia is sometimes prepared from the residues of the beet sugar 
industry, from the refuse of slaughter houses and tanneries, and from the 
gases from coke ovens. It is not obtained directly from the nitrogen of 
the air. 

Properties of Ammonia. Ammonia gas is colorless. 
It has an exceedingly pungent odor, and if inhaled sud- 
denly or in large quantities it brings tears to the eyes and 
may cause suffocation. It is a light, volatile gas, being only 
.59 times as heavy as air. A liter of the gas at o and 
760 mm. weighs .77 gm. It will not burn in the air, nor 
will it support the combustion of a blazing stick ; but if the 
air is heated or if its proportion of oxygen is increased, a 
jet of ammonia gas will burn in it with a yellowish flame, 
thereby illustrating the. broader application of the term 

Ammonia gas is easily liquefied if reduced to oC. and 
subjected to a pressure of 4^ atmospheres, while at 34 C. 
it liquefies at the ordinary atmospheric pressure. 

Compounds of Nitrogen. 149 

Liquefied ammonia is often called anhydrous ammonia, because it 
contains no water. It boils at 33. 5 C. Hence, if it is exposed to 
the air or warmed in any way, it changes back to a gas, and in so doing 
absorbs considerable heat. This fact has led to the extensive use of 
liquid ammonia in the manufacture of ice. 

Ammonia is a strong alkali, and was called formerly the 
volatile alkali. Priestley, who discovered and studied the 
gas, called it alkaline air. 

Another marked property of ammonia gas is its solu- 
bility in water. A liter of water at oC. dissolves 1148!. 
of gas (measured at OC. and 760 mm.), and at the 
ordinary temperature I 1. of water dissolves about 700 1. of 
gas. This solution of the gas is usually called ammonia, 
though other names, especially ammonium hydroxide, are 
sometimes applied to it. Commercially it is known as 
aqua ammonia, ammonia, or ammonia water. It gives off 
the gas freely, when heated, as may easily be discovered 
by the odor or by the formation of the dense white fumes 
of ammonium chloride (NH 4 C1) when the solution is ex- 
posed to hydrochloric acid. The solution is lighter than 
water, its specific gravity being about .88, and contains 
about 35 per cent (by weight) of the gas. It is a strong 
alkali a caustic alkali, neutralizes acids and forms salts, 
and acts in many respects like sodium hydroxide. 

Ammonium Hydroxide and Ammonium Compounds. 

When ammonia gas is passed into water^it is believed that 
the ammonia combines with the water and forms a solution 
of an unstable compound having the formula NH 4 OH. 
This compound is ammonium hydroxide (or ammonium 
hydrate). Its formation may be represented thus 

NH 3 -f- H 2 O NH 4 OH 

Ammonia Water Ammonium Hydroxide 

150 Descriptive Chemistry. 

Ammonium -hydroxide acts like a base. It has a marked 
alkaline reaction ; it neutralizes acids and forms salts, 

NH 4 OH + HC1 NH 4 C1 + H 2 O 

Ammonium Chloride 

2NH 4 OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2 H 2 O 

Ammonium Sulphate 

These salts, ammonium chloride and ammonium sul- 
phate, have definite properties, and are strictly analogous 
to sodium salts. Thus, we have 

Sodium Salts Ammonium Salts 
Nad NH 4 C1 

NaNO 3 * NH 4 NO 3 

Na 2 S0 4 (NH 4 ) 2 S0 4 

etc. etc. 

Hence, it is believed that ammonium compounds contain a 
group of atoms which acts like an atom of a metal. This 
group of atoms is called ammonium, and its formula is 
NH 4 . Ammonium has never been separated from its 
compounds, or if it has it is so unstable that it immedi- 
ately decomposes into ammonia gas and hydrogen. So 
also ammonium hydroxide has never been obtained free, 
for it decomposes readily into ammonia gas and water, 


NH 4 OH NH 3 + H 2 O 

Ammonium Hydroxide Ammonia Gas Water 

Ammonium is sometimes called a radical, because it is the root or 
foundation of a series of compounds. It is likewise called a hypotheti- 
cal metal, because its existence is assumed and it acts chemically like 

Compounds of Nitrogen. 151 

Ammonium Chloride is prepared by passing ammonia 
gas into dilute hydrochloric acid, by mixing ammonium 
hydroxide and hydrochloric acid, or by letting the two 
gases mingle. The equation for the essential reaction is 

NH 3 + HC1 = NH 4 C1 

Ammonia Hydrochloric Acid Ammonium Chloride 

It is convenient to regard this compound as the ammonium 
salt of hydrochloric acid, as if it were formed by replacing 
the hydrogen of the acid by ammonium, just as sodium 
forms sodium chloride. 

Ammonium chloride is a white, granular or crystalline 
solid, with a sharp, salty taste. It dissolves easily in 
water, and in so doing lowers the temperature markedly. 
When heated to a high temperature it gradually breaks 
up into ammonia and hydrochloric acid. This kind of 
decomposition is called dissociation. 

Large quantities of ammonium chloride are made at one stage of the 
manufacture of ammonium hydroxide by passing the gas into hydro- 
chloric acid. The crude product is called " muriate of ammonia " to 
indicate its relation to muriatic (or hydrochloric) acid. It is largely 
used for charging Leclanche' batteries, as an ingredient of soldering 
fluids, in galvanizing iron, and in textile industries. The crude salt is 
purified by heating it gently in a large iron or earthenware pot, with a 
dome-shaped cover ; the ammonium chloride volatilizes easily and then 
crystallizes in the pure state as a fibrous mass on the inside of the cover, 
but the impurities remain behind in the vessel. The process of vapor- 
izing a solid substance and then condensing the vapor directly into 
the solid state is called sublimation. It differs from distillation in that 
the substance does not pass through an intermediate liquid state. The 
product of sublimation is called a sublimate. Sublimed ammonium 
chloride is known as sal ammoniac. 

Ammonium Sulphate is made by passing ammonia gas into sul- 
phuric acid, or by adding ammonium hydroxide to the acid, thus 

2NH 4 OH + H 2 SO 4 = (NH 4 ) 2 SO 4 + 2 H 2 O 
Ammonium Hydroxide Ammonium Sulphate 

152 Descriptive Chemistry. 

* The commercial salt is a grayish or yellowish solid. It is used as a 
Constituent of fertilizers, since it is rich in nitrogen, and in making 
ammonium alum and other ammonium compounds. 

Ammonium Nitrate is made by passing ammonia into nitric acid, or 
by allowing ammonia gas and the vapor of nitric acid to mingle, thus 

NH 3 + HNO 3 = NH 4 NO 3 

Ammonia Nitric Acid Ammonium Nitrate 

It is a white salt which forms beautiful crystals. It dissolves easily in 
water with. a fall of temperature. Its chief use is in the preparation of 
nitrous oxide (see this compound). 

Ammonium Carbonate is an impure salt as found in commerce, 
being a mixture of acid ammonium carbonate (HNH 4 CO 3 ) and a 
related compound. When pure and fresh it is transparent, but on ex- 
posure to the air it loses ammonia and turns white. It is used to pre- 
pare some kinds of baking powder, to scour wool, as a medicine, and 
to prepare smelling salts, since it gives off ammonia readily. 

Other ammonium compounds are sodium ammonium phosphate 
or microcosmic salt (HNaNH 4 PO 4 ), ammonium sulphocyanate 
(NH 4 SCN), and ammonium sulphide ( (NH 4 ) 2 S). 

Uses of Ammonia. Ammonia in the different forms is 
widely used as a cleansing agent, especially for the re- 
moval of grease, as a restorative in cases of fainting or of 
inhaling irritating gases, in dyeing and calico printing, and 
in the manufacture of dyestuffs, sodium carbonate, and 
ice. Its salts have many domestic, industrial, and agri- 
cultural uses. 

The Use of Ammonia as a Refrigerant and in making 
Ice depends upon the fact that many liquids in passing 
into a gas absorb heat. Liquefied ammonia (not the ordi- 
nary liquid ammonia) changes rapidly into a gas when its 
temperature is raised or the pressure reduced. Hence, if 
anhydrous ammonia is allowed to flow through a pipe sur- 
rounded by brine, the ammonia evaporates in the pipe and 
cools the brine, which may be used as a refrigerant or for 

Compounds of Nitrogen. 153 

making ice. In some cold storage houses, breweries, 
packing houses, and sugar refineries, this cold brine is 
pumped through pipes placed in the rooms where a low 
temperature is desired. 

The construction and operation of an ice-making plant are essentially 
as follows : 

Liquefied ammonia is forced from a tank into a series of pipes which 
are submerged in an immense vat filled with brine. Large galvanized 
iron cans containing pure water to be frozen are immersed in the brine, 
which is being kept below the freezing point of water by the rapid evap- 
oration of the ammonia in the pipes. In about sixty hours the water 
in the cans is changed into a cake of ice weighing about three hundred 
pounds. As fast as the ammonia gas forms in the pipes, it is removed 
by exhaust pumps into another tank, where it is recondensed to liquefied 
ammonia and conducted, as needed, into the first tank to be used again. 
The ammonia is thus used over and over without appreciable loss. 
The pure water is sometimes obtained by condensing the exhaust 
steam from the boilers used to operate the machinery, though it 
usually comes from a deep well. Most ocean steamers have an ice 
plant, and in large cities in warm climates manufactured ice is a com- 
mon commodity. 

Composition of Ammonia Gas. Numerous experiments show that 
ammonia gas has the composition expressed by the formula NH 3 . 

(1) Dry ammonia gas passed over heated magnesium decomposes into 
hydrogen and nitrogen. The hydrogen may be collected and tested, 
but the nitrogen combines with the magnesium, forming a yellowish 
green powder called magnesium nitride, thus 

2NH 3 + 3 Mg = Mg,N 2 + 3 H 2 

Magnesium Magnesium Nitride 

These facts show that ammonia contains nitrogen and hydrogen. 

(2) If a bottle is filled with chlorine gas and plunged mouth downward 
into a vessel containing ammonium hydroxide, dense white fumes fill 
the bottle, the greenish chlorine gas disappears, and the liquid rises in 
the bottle ; after the bottle has stood mouth downward in a dish con- 
taining dilute hydrochloric acid (to neutralize the excess of ammonia), 
the gas in the bottle will be found to be nitrogen. The chlorine with- 

154 Descriptive Chemistry. 

draws the hydrogen from the ammonia of the ammonium hydroxide, 
leaving the nitrogen free, thus 

NH, + 3 Q - N + 3 HC1 

Ammonia Chlorine Nitrogen Hydrochloric Acid 

(3) The same experiment, if performed accurately, shows that one 
volume of nitrogen combines with three volumes of hydrogen to form 
ammonia gas. A tube containing a known volume of chlorine is pro- 
vided with a funnel through which concentrated ammonium hydroxide 
is dropped into the chlorine, until the reaction ceases (Fig. 22). After 
the excess of ammonia is neutralized with sulphuric acid, the volume 
of nitrogen left is one third of the original volume 
of chlorine gas. Now hydrogen and chlorine com- 
bine in equal volumes, hence the volume of hydrogen 
withdrawn from the added ammonia must be equal 
to the original volume of chlorine. But this volume 
is three times the volume of nitrogen, therefore there 
must be three times as much hydrogen as nitrogen 
in ammonia gas. (4) When electric sparks are 
passed through ammonia gas, it is decomposed into 
nitrogen and hydrogen. Now if oxygen is added, 
and an electric spark passed through the mixture, the 
oxygen and hydrogen combine. The volume of the 
remaining nitrogen is one fourth of the mixture of 
nitrogen and hydrogen, hence the hydrogen must 
have been three fourths ; that is, the volume of 
FIG. 22. Appa- hydrogen in the original volume ammonia was three 
ratus for determin- t i mes t i iat o f the nitrogen. (5) The gravimetric 
iner the composition ... r . c , , .,. . 

jr composition or ammonia gas is found by oxidizing 

it, and weighing the water and nitrogen, which are 
the only products. The result shows that fourteen parts of nitrogen 
combine with three parts of hydrogen. (6) The vapor density has 
been found to be 8.5. These facts require NH 3 as the simplest formula 
for ammonia and 17 as its molecular weight. Independent experiments 
verify this molecular weight. 


Nitric Acid is one of the most useful compounds of 
nitrogen. It was known to the alchemists, who used it 

Compounds of Nitrogen. 155 

to prepare a mixture which dissolves gold. Nitric acid 
is used in the preparation of many nitrogen compounds. 

Formation of Nitric Acid. When moist animal or 
vegetable matter containing nitrogen decays in the presence 
of an alkali, nitric acid is formed ; it is neutralized at once 
by the alkali, so nitrates salts of nitric acid are the 
final products. This chemical change is known as nitri- 
fication, and it is caused, or largely influenced, by minute 
living organisms called bacteria. Nitrification is constantly 
going on in the soil and is an exceedingly helpful process, 
since it transforms harmful waste matter into valuable 
plant food. 

As a result of nitrification, there are vast deposits of nitrates, espe- 
cially in desert regions and tropical countries. For example, potassium 
nitrate (KNO 3 ) is found in the soils near large cities in India, Persia, 
and Egypt. 

Nitric acid is formed in small quantities when electric 
sparks are passed through moist air. Hence nitric acid or 
its salts can be detected in the atmosphere after a thunder- 

This chemical change is now being applied on a large scale at Ni- 
agara Falls. Electric sparks are passed through confined air and the 
products are forced into a tower. Here they are absorbed in water or 
in a solution of lime ; thereby forming nitric acid or calcium nitrate. 
The latter is converted into sodium nitrate (see below). 

Preparation. Nitric acid is prepared in the laboratory 
by heating concentrated sulphuric acid with a nitrate, usu- 
ally sodium or potassium nitrate. About equal weights of 
nitrate and acid are put into a glass retort and gently 
heated. The nitric acid distils into a receiver, which is 
kept cool by running water, ice, or moist paper. The 

Descriptive Chemistry. 

chemical change at a low temperature is represented by 
the equation 

NaNO 3 + H 2 SO 4 = HNO 3 + HNaSO 4 
Sodium Nitrate Sulphuric Acid Nitric Acid Acid Sodium Sulphate 
85 +98 = 63 + 120 

But if the temperature is high and an excess of the nitrate 
is present, the equation is 


H 2 SO 4 = 2HNO 3 

Na 2 SO 4 
170 +98 126 4- 142 

A high temperature, however, decomposes part of the 
nitric acid, hence excessive heat is usually avoided. 

FlG. 23. Apparatus for the manufacture of nitric acid. 

Nitric acid is manufactured on a large scale by heating sodium nitrate 
and sulphuric acid in a large cast-iron retort (A) connected with huge 
glass or earthenware bottles (Z?, B, ), arranged as shown in Figure 23 ; 
the last bottle is connected with a tower filled with coke over which 
water trickles to absorb the vapors which escape from the bottles. The 
acid vapors are also often absorbed in earthenware or glass tubes. 

Properties. Pure nitric acid is a colorless liquid, but 
the commercial acid is yellow or reddish, due to absorbed 
nitrogen compounds, chlorine, or iron compounds. It de- 
composes slowly in the sunlight or when heated, and a 

Compounds of Nitrogen. 157 

brownish gas may often be seen in bottles of nitric acid. 
It absorbs water, and forms irritating fumes when exposed 
to the air. The specific gravity of the commercial acid is 
about 1.42, and it contains from 60 to 70 per cent of the 
real acid (HNO 3 ), the rest being water. 

If the water is removed by slowly distilling the commercial acid with 
concentrated sulphuric acid, the product contains from 94 to 99 per cent 
of the real acid and its specific gravity is about 1.51. When nitric 
acid is boiled, it loses either acid or water until the liquid contains 
approximately 68 per cent of nitric acid, and then it continues to 
boil unchanged at 120 C. 

Nitric acid is very corrosive. It turns the skin a perma- 
nent yellow color, and may cause serious burns. Many 
organic substances are turned yellow and sometimes com- 
pletely decomposed by it. It parts readily with its oxygen, 
especially when hot, and is therefore an energetic oxidizing 
agent. Charcoal burns brilliantly in hot acid, while straw, 
sawdust, hair, and similar substances are charred and even 
inflamed by it. Iron sulphide heated with nitric acid 
becomes iron sulphate, by the addition of oxygen, thus 

FeS + 2O 2 = FeSO 4 

Iron Sulphide Oxygen Iron Sulphate 

Uses of Nitric Acid. Nitric acid is one of the com- 
mon laboratory acids. Large quantities are used in the 
manufacture of nitrates, dyestuffs, sulphuric acid, nitro- 
glycerine, gun cotton, in the refining of gold and silver, 
and in etching copper plates. 

Composition of Nitric Acid. Although the alchemists knew and 
valued nitric acid, its composition was a mystery until Lavoisier showed 
in 1776 that it contained oxygen and probably nitrogen. Its exact 
composition was determined by Cavendish in 1784-1785, by passing 
electric sparks through a mixture of oxygen and nitrogen in the pres- 


Descriptive Chemistry. 

ence of water or caustic potash. The same facts had been observed, but 
not explained, by Priestley. Many independent experiments show that 
the composition of nitric acid is expressed by the formula HNO 3 . 

(1) When electric sparks are passed through a bottle containing moist 
air or a solution of potassium hydroxide, the water becomes acid to 
litmus or the liquid will be found to contain a trace of potassium nitrate. 

(2) Nitric acid may be reduced to ammonia by nascent hydrogen, thus 
showing that the acid contains nitrogen. (3) Conversely, if a mixture 
of ammonia and air is passed over a mass of hot, porous platinum, 
nitric acid is formed. (4) If the acid is allowed to flow through a hot 
porcelain or clay tube, oxygen is one of the gaseous products. 

Nitrates. Nitric acid is monobasic and forms a series 
of well-defined salts called nitrates. The interaction of 
nitric acid and most metals is exceedingly vigorous, and 
for this reason, probably, the alchemists called the acid 
aquafortis strong water. The reaction varies with the 
metal, strength of the acid, temperature, and the presence 
of resulting compounds. 

The solid product of the reaction is usually a nitrate, though some 
metals, such as tin and antimony, form oxides. The gaseous products 
are usually oxides of nitrogen, especially nitric oxide (NO), which, 
however, quickly forms nitrogen peroxide (NO 2 ) in the air. Hydrogen 
is never liberated so that it can be collected ; probably it immediately 
reduces the nitric acid to another compound of nitrogen. Nitrates are 
also formed by the action of nitric acid upon oxides, hydroxides, and 
carbonates, thus 

Copper Oxide 



Na 2 CO 3 



2HNO = 

HN0 - 

Cu(NO,) 2 
Copper Nitrate 

KN0 3 


2HNO 3 = 2 NaNO 3 


H 2 

H 2 O 

H 2 O 

Compounds of Nitrogen. 159 

When nitric acid is poured upon copper, the liquid bub- 
bles violently and becomes hot, dense fumes of a reddish 
brown gas are given off, and the liquid turns blue owing 
to the dissolved copper nitrate. Other metals, such as 
zinc, iron, and silver, act in a similar way, though the nitrate 
is blue only in the case of copper. The usual equation for 
the chemical change with copper is 

3Cu + 8HNO 3 = 3Cu(NO 3 ) 2 + 2 NO + 4 H 2 O 
Copper Nitrate Nitric Oxide 

When nitric oxide is exposed to the air, it changes at once 
into the reddish brown peroxide, thus 

NO + O NO 2 

Nitric Oxide Oxygen Nitrogen Peroxide 

Nitrates as a rule are very soluble in water. They be- 
have in various ways when heated. Some, like sodium 
and potassium nitrates, lose oxygen and pass into nitrites ; 
others, like copper nitrate, form an oxide of the metal, an 
oxide of nitrogen, and oxygen ; and one, ammonium 
nitrate, decomposes into water and nitrous oxide (N 2 O). 
Since many nitrates, when heated, give up oxygen, they 
are powerful oxidizing agents. Potassium nitrate dropped 
on hot charcoal burns the charcoal vigorously and rapidly. 
This kind of chemical action is called deflagration. 

The Test for Nitrates (and of course for nitric acid) is as follows : 
Add to the solution of the nitrate a little concentrated sulphuric acid, 
and upon the cool mixture pour carefully a cold, dilute solution of fresh 
ferrous sulphate. A brown layer is formed where the two liquids meet. 

Nitrous Acid (HNO 2 ) has never been obtained in the free state, but 
its salts the nitrites are well known. Potassium nitrite (KNO 2 ) 
and sodium nitrite (NaNO 2 ) are formed by removing the oxygen from 
the corresponding nitrate by heating gently or by heating with lead. 
Nitrites give off brown fumes when treated with sulphuric acid, thus 


Descriptive Chemistry. 

being readily di 
decomposition o 
amount in drin 

shed from nitrates. Nitrites are formed by the 
' C matter, and the presence of a relatively large 
r indicates contamination by sewage. 

Aqua Regia is an old term which is still applied to a 
mixture of concentrated nitric and hydrochloric acids. 
The expression means "royal water," and indicates that 
the mixture dissolves gold and platinum the noble metals. 
Its solvent power depends mainly upon the free chlorine 
which is produced in the mixture by the oxidizing action 
of the nitric acid. The product of the action of aqua 
regia on metals is always the chloride of the metal. 

Oxides of Nitrogen. There are five oxides of ni- 
trogen : 




Nitrous oxide 

N 2 O 

Colorless "as 


Colorless o'as 

Nitrogen trioxide 
Nitrogen peroxide 


NO 9 

Blue liquid 
Brown gas 

Nitrogen pentoxide . . 


White solid 


Only three of these are important, viz., nitrous and nitric 
oxides, and nitrogen peroxide. 

Nitrous Oxide is one of the numerous decomposition 
products of nitric acid, but it is usually prepared by decom- 
posing ammonium nitrate. This salt, if gently heated in a 
test tube provided with a delivery tube, first melts and then 
decomposes into water and nitrous oxide ; the gas may be 
collected over warm water. The equation of the chemical 
change is 

NH 4 NO 3 = N 2 O 

Ammonium Nitrate Nitrous Oxide 

2H 2 O 

Compounds of Nitrogen. 161 

This colorless gas has a sweet taste and a faint but pleas- 
ant odor. It is less soluble in hot than in cold water. The 
gas does not burn, but it supports the combustion of many 
burning substances, though not so vigorously as oxygen 
does. Sulphur, for example, will not burn in nitrous oxide, 
unless the sulphur is hot and well ignited at first. The 
most striking property of nitrous oxide is its effect on the 
human system. If breathed for a short time, it causes 
more or less nervous excitement, often manifested by 
laughter, and on this account the gas was called "laughing 
gas" by Davy. If breathed in large quantities, it slowly pro- 
duces unconsciousness and insensibility to pain. The gas 
is often used when insensibility is desired for a short time, 
as in dentistry. 

It is easily liquefied by cold and pressure, and is often used in this 
form to furnish the gas itself and to produce very low temperatures. It 
is a commercial article and is sold in small iron cylinders. 

Nitrous oxide was discovered by Priestley in 1776; but its composi- 
tion was not explained until 1799, when Davy, by an extensive study of 
its properties, proved it to be an oxide of nitrogen. In his enthusiasm 
Davy wrote a friend: "This gas raised my pulse upward of twenty 
strokes, made me dance about the laboratory as a madman, and has 
kept my spirits in a glow ever since." It is needless to say that the 
usual results are more quieting. 

The Composition of Nitrous Oxide is shown as follows : By ex- 
oloding equal volumes of nitrous oxide and hydrogen, only nitrogen 
Remains, and its volume equals the original volume of nitrous oxide. 
The oxygen unites with the hydrogen to form water, and there is just 
enough oxygen to unite with a volume of hydrogen equal to the volume 
of the nitrous oxide. Therefore, the oxygen in the nitrous oxide must 
have been equal to half the volume of the nitrogen, since oxygen and 
hydrogen combine in the ratio of one to two. Furthermore, experiment 
has shown that the weights of equal volumes of nitrous oxide and ni- 
trogen are in the ratio of 44 to 28. Therefore, the smallest part of 
oxygen united with the nitrogen must weigh 16 ; and since the nitrogen 
weighs 28, the formula must be N.,0. 

1 62 Descriptive Chemistry. 

Nitric Oxide has long been known, since it is the usual 
gaseous product of the interaction of nitric acid and metals. 
It is usually prepared by the interaction of copper and 
dilute nitric acid (sp. gr. 1.2). The equation for the com- 
plex chemical change is usually written thus 

3Cu + 8HNO 3 = 2 NO + Cu(NO 3 ) 2 + 4 H 2 O 
Copper Nitric Acid Nitric Oxide Copper Nitrate 

The gas thus prepared is impure, and it is customary to 
use ferrous sulphate and nitric acid as a source of the 
pure gas. 

Nitric oxide is a colorless gas, but upon exposure to the 
air, it combines at once with oxygen, forming dense red- 
dish brown fumes of hydrogen peroxide. The simplest 
equation for this change is 

NO + O = NO 2 

Nitric Oxide Nitrogen Peroxide 

This property distinguishes nitric oxide from all other 
gases. It does not burn, nor does it support combustion 
unless the burning substance (e.g. phosphorus or sodium) 
introduced is hot enough to decompose the gas into nitro- 
gen and oxygen, and then, of course, the liberated oxygen 
assists the combustion. 

The Composition of Nitric Oxide is determined by heating iron or 
another metal in it. The oxygen of the oxide combines with the iron, 
and the nitrogen is left free. The resulting volume of nitrogen is half 
the volume of the nitric oxide taken. Hence nitric oxide contains 
equal volumes of nitrogen and oxygen. By an independent experiment 
the molecular weight is found to be 30. Hence the formula must be NO. 

Nitrogen Peroxide is the reddish brown gas formed by 
the direct combination of nitric oxide and oxygen. Thus 

NO + O NO 2 

Nitric Oxide Nitrogen Peroxide 

Compounds of Nitrogen. 163 

It is also produced by heating certain nitrates. Thus 

Pb(NO 3 ) 2 = 2NO 2 + PbO + O 
Lead Nitrate Nitric Oxide Lead Oxide Oxygen 

The fumes of nitrogen peroxide always appear when nitric 
acid and metals interact, but, as already stated, the fumes 
are not produced at first, being the result of a second 
chemical change when the real product, nitric oxide, 
comes in contact with oxygen of the air. 

Nitrogen peroxide is poisonous. It dissolves in water ; 
it also dissolves in concentrated nitric acid, forming 
fuming nitric acid. 

At very low temperatures nitrogen peroxide is a colorless solid. At 
about 10 C. it is a yellowish liquid, and as the temperature rises the 
color grows darker, until at 22 C. the liquid boils and gives off the 
familiar reddish brown gas. Above 140 C. this gas begins to lose its 
color, and at 600 C. the color entirely disappears. The density of the 
gas at low temperatures indicates the formula N 2 O 4 , whence the name 
nitrogen tetroxide, often used. But the density at about 140 C. 
indicates the formula NO 2 . 

Nitrogen Trioxide, N 2 O 3 , and Nitrogen Pentoxide, N 2 O 5 , are unstable 
compounds and have no practical importance. They are the anhy- 
drides of nitrous and nitric acids, thus 

N 2 8 + H 2 O 2HNO 2 

Nitrogen Trioxide Nitrous Acid 

N 2 5 + H 2 O 2HNO ? 

Nitrogen Pentoxide Nitric Acid 


1. Name several sources of ammonia gas. How is ammonia gas 
prepared in the ' laboratory ? Give the equation for the reaction. 
State its important properties. 

2. What is ammonium hydroxide ? How is it prepared on a large 
scale ? Summarize its properties. What are its uses ? 

3. What is the meaning and significance of (a) volatile alkali, 

164 Descriptive Chemistry. 

($) anhydrous ammonia, (V) spirits of hartshorn, (d) sal volatile, 
(^) muriate of ammonia, (/") sal ammoniac, (g) aqua for'tis ? 

4. Why is NH 3 the formula of ammonia gas ? 

5. Give several tests for (a) ammonia, and (V) nitric acid. 

6. What different meanings may the word ammonia have ? What 
is ammoniacal liquor? Gas liquor? Aqua ammonia? Ammonium 
hydrate ? Ammonia of commerce ? Ammonia water ? 

7. How is ammonia gas liquefied ? Describe the manufacture of 
ice by liquid ammonia. 

8. Develop the topics : (a) ammonium is a radical ; () nitric acid 
is an oxidizing agent ; (<:) nitrates are unstable ; (d) fuming nitric acid. 

9. Give the formula, method of preparation, properties, and uses of 
(a) ammonium chloride, (b) ammonium nitrate, (c} ammonium sulphate, 
(d) ammonium carbonate. 

10. How is nitric acid formed (a) in the soil, () in the air ? How 
is it prepared (a) in the laboratory, (b) on a large scale ? Summarize 
(a) the physical properties of nitric acid, and (b} its chemical properties. 
For what is it used ? 

1 1 . What is the formula of nitric acid ? Summarize the evidence 
of its composition. 

12. What are nitrates ? How are they formed ? What is the 
effect of heat upon () potassium nitrate, ($) copper nitrate, (c} am- 
monium nitrate ? Give other properties of nitrates. What is the test 
for nitrates ? 

13. What are nitrites ? How are they formed ? How are they 
distinguished from nitrates ? 

14. What is aqua regia? For what is it used ? Why so called ? 
What is the chemical action of aqua regia on gold ? Upon what prop- 
erty of nitric acid does its chemical action depend ? 

15. Give the names and formulas of the five oxides of nitrogen. 
Describe the preparation of nitrous oxide. State briefly its properties. 
For what is it used ? Who discovered it ? What did Davy call it ? 
Why ? Summarize the evidence of the composition of nitrous oxide. 

1 6. Describe the preparation of nitric oxide. State the equation 
for the reaction. What are its properties ? 

17. How is nitrogen peroxide prepared ? State its properties. 
How is it readily distinguished from all other oxides of nitrogen ? 
What two formulas have been given to nitrogen peroxide ? Why ? 

1 8. What is (#) nitric oxide, (b) nitrous oxide, (c) nitrogen per- 

Compounds of Nitrogen. 165 

oxide, (//) nitrogen tetroxide, (c) nitrogen trioxide, (d) nitrogen 
monoxide, (e) nitrogen pentoxide ? 

19. State the equation for the preparation of (a) nitric acid at a 
low temperature, () nitric acid at a high temperature, (V) ammonium 
chloride, (d) ammonium hydroxide from water and ammonia, (d} ni- 
trous oxide, (e) nitrogen peroxide, (/") copper nitrate. 

20. Define and illustrate () sublimation, ($) sublimate, (V) nitrifi- 
cation, (d} deflagration, (>) nitrate, (/") ammonium compound. 

21. What is the valence of nitrogen in ammonia gas ? In ammo- 
nium ? In ammonium hydroxide ? 

22. (a) Why are there no acid nitrates ? () What is the valence 
of nitrogen in nitric acid, copper nitrate, nitrous oxide, nitric oxide, 
nitrogen peroxide, nitrogen trioxide, nitrogen pentoxide ? 


1. How many grams of ammonia gas can be obtained from 2140 
gm. of ammonium chloride by heating with lime ? 

2. Calculate the percentage composition of (a) ammonium chloride, 
() ammonium hydroxide, (Y) ammonium sulphate, (d} ammonium 

3. Calculate the simplest formula of the compounds having the per- 
centage composition (a) N = 82.35, H = 17.64; and () N = 26.17, 
Cl = 66.35, H = 7.48. 

4. Calculate the percentage composition of (a) nitric acid, () po- 
tassium nitrate (KNO 3 ), (V) sodium nitrate. 

5 . How many grams of nitric acid can be obtained by heating a 
kilogram of sodium nitrate with sulphuric acid at a low temperature ? 

6. If the specific gravity of a sample of nitric acid is 1.522, 
(a) what will 100 cc. weigh, and (b) what volume must be taken to 
weigh 100 grams ? 

7. Calculate the simplest formula of the substances having the 
composition (a) O = 76.19, H = 1.58, N = 22.22; () N =13.86, 
K = 38.61, O = 47.52. 



Properties of Gases. Extensive study of gases shows 
that they all conform to simple laws. Thus we have 
already seen that they behave uniformly with changes of 
pressure (Boyle's law) and with changes of temperature 
(Charles's law). Other simple relations prevail. 

Gay-Lussac's Law. Gases combine by volume in 
simple ratios. Experiment has revealed the following 
facts about the 




2 vol. hydrogen 
I vol. oxygen 

2 vol. water vapor 

I vol. chlorine 
i vol. hydrogen 

2 vol. hydrochloric acid gas 

3 vol. hydrogen 
i vol. nitrogen 

2 vol. ammonia gas 

2 vol. nitrogen 
i vol. oxygen 

2 vol. nitrous oxide gas 

2 vol. nitrogen 
3 vol. oxygen 

2 vol. nitrogen trioxide gas 

1 66 

Avogadro's Hypothesis. 167 

Additional illustrations will be given in later chapters. The 
simple ratio which exists between the gas volumes, whether 
components or products, has been found to be true of all 
gases. The law was pointed out in 1808 by Gay-Lussac, 
who stated the relation substantially as follows: 

Gases combine in volumes which bear a simple ratio to 
each other and to that of the product. 

By " a simple ratio " we mean one made up of small 
whole numbers. As a rule, the product occupies two unit 

"" Avogadro's Hypothesis. In 1811 an Italian physi- 
cist proposed an hypothesis to account for the similar 
behavior of gases. At that time the properties of gases 
were not generally known, and the views of Avogadro 
were overlooked until about 1860. Since then the hypo- 
thesis has been helpful in explaining many facts, and it 
is generally accepted by chemists as a very probable 
assumption. It may be stated thus: 

There is an equal number of molecules in equal vohimes 
of all gases at the same temperature and pressure. 

This statement cannot be proved directly by experiment, 
but there is much physical, chemical, and mathematical 
evidence in harmony with it. 

According to Avogadro's hypothesis a liter of hydrogen 
and a liter of oxygen at the same temperature and pres- 
sure contain the same number of molecules, though we 
do not know how many. Suppose, however, that each 
liter contained 1000 molecules. A liter of hydrogen 
weighs 0.0896 gm. and a liter of oxygen at the same tem- 
perature and pressure weighs 1.43 gm. But 0.0896 and 
1.43 are in the same ratio as i and 16. Therefore, since 

1 68 Descriptive Chemistry. 

a thousand molecules of oxygen weighs 16 times more than 
a thousand molecules of hydrogen, a single molecule of 
oxygen must weigh 16 times more than a single molecule 
of hydrogen. Therefore, in general, in order to find how 
much heavier any gaseous molecule is than a hydrogen 
molecule, it is only necessary to compare the weights of 
equal volumes of hydrogen and the gas under examination. 

An application of Avogadro's hypothesis is made in course of the 
following argument, which proves that a molecule of hydrogen consists 
of two atoms : 

One volume of hydrogen combines with one volume of chlorine to 
form two volumes of hydrochloric acid gas. Suppose the volume of 
hydrogen contained 100 molecules. Then, according to Avogadro's 
hypothesis, the equal volume of chlorine will contain 100 molecules, 
while the two volumes of the product will contain 200 molecules of 
hydrochloric acid gas. That is 

100 molecules of Hydrogen + 100 molecules of Chlorine 
= 200 molecules of Hydrochloric Acid Gas. 

Now every molecule of hydrochloric acid gas contains at least one atom 
each of hydrogen and chlorine, and the 200 molecules must contain 
200 atoms each of chlorine and hydrogen. Therefore each molecule of 
hydrogen and of chlorine must be divisible into two atoms, since the 
100 hydrogen and the 100 chlorine molecules provide the 200 hydrogen 
atoms and the 200 chlorine atoms in the 200 molecules of hydrochloric 
acid gas. Similar reasoning leads to the conclusion that the molecules 
of oxygen, nitrogen, and most elementary gases consist of two atoms. 

Vapor Density and Molecular Weight. It was stated 
in a previous chapter that a molecular weight is the sum 
of the weights of the atoms in the molecule. But this 
method of finding the molecular weight is useless, unless 
we first know the formula, and in many cases the formula 
cannot be chosen until after the molecular weigjit has been 
found by several methods. Hence, the determination of 
molecular weights is an important matter. In the case 

Vapor Density and Molecular Weight. 169 

of gaseous or volatile elements and compounds, it is often 
accomplished by finding the vapor density of the substance. 
There is a direct and simple relation between molecular 
weight and vapor density. By vapor density we mean the 
ratio of the weight of a gas to the weight of an equal 
volume of hydrogen at the same temperature and pressure. 
Thus, the vapor density of steam is 9, because experiment 
shows that it weighs 9 times more than an equal volume of 
hydrogen under the same conditions of temperature and 
pressure. Therefore the molecular weight of steam is 9 
times the molecular" weight of hydrogen. But the molec- 
ular weight of hydrogen is 2, since its molecule contains 
two atoms each weighing I. Therefore, the molecular 
weight of steam is 18, or twice the vapor density. The 
general fact that the molecular weight of a gaseous com- 
pound is twice its vapor density is clearly seen from the 
following table showing the 





Carbon dioxide . 
Ammonia. .... 



Hydrochloric acid 
Water vapor (steam) . 





Hence, a determination of the vapor density of a com- 
pound or an element allows us to select the correct molec- 
ular weight and assign the proper formula. 

The vapor densities of the elements mercury and cadmium show 
that the atom and molecule are identical, while the vapor densities of 
phosphorus and arsenic indicate that the molecule of each consists of 
four atoms. A molecule of oxygen contains two atoms, but a molecule 
of ozone contains three ; therefore, the formula of ozone is O 3 . 

i jo Descriptive Chemistry. 

Other Methods of determining Molecular Weights. Some sub- 
stances cannot be vaporized without decomposition. The molecular 
weights of such substances cannot, of course, be found by the vapor 
density method. If a substance dissolves without decomposition, its 
molecular weight can be determined by the boiling-point or freezing- 
point method, which was briefly described in Chapter X. The above 
methods give approximate results. Exact molecular weights are found 
by accurate quantitative analysis. Suppose we wished to find the molec- 
ular weight of acetic acid. Silver acetate is analyzed and found to con- 
tain 64.65 per cent of silver ; the per cent of the remaining elements 
of the molecule must be 35 35. The atomic weight of silver is 107.93, 
if the atomic weight of oxygen is 16. Hence, the weight of the silver 
acetate molecule, except the silver, is found by the proportion 

107.93: x : : 64.65 : 35.35. x= 59.02. 

Silver acetate is formed by replacing one atom of the hydrogen of the 
acid by one atom of silver. Therefore, the weight of the molecule of 
acetic acid is found by adding to 59.02 the weight of one atom of hydro- 
gen. That is, the exact molecular weight of acetic acid is 60.028 
(i.e. 59.02 + 1.008). 

Determination of Atomic Weights. The atomic 
weight of an element, as already stated, is a relative weight. 
It is a number expressing the relation of the weight 
of an atom of a given element to the weight of an atom 
of some element chosen as a standard. Thus, if we say 
that the atomic weight of nitrogen is 14, we mean that 
the relation between the weight of the nitrogen atom and 
that of the hydrogen atom is 14 to i, if we adopt the 
hydrogen atom as the standard atom ; or we mean that 
the relation between the weight of the nitrogen atom and 
that of the oxygen atom is 14 to 16, if we adopt the oxygen 
atom as the standard. The approximate atomic weights 
are usually expressed in round numbers, and do not 
vary much with the standard. Wherever exact atomic 
weights are used in this book, the oxygen standard is the 

Determination of Atomic Weights. 171 

In Chapter IX it was stated that the determination and 
selection of atomic weights are based on several principles. 
This subject can now be appropriately considered. 

One method of selecting the atomic weight is illustrated 
by the case of chlorine, which has the atomic weight 35.5. 
The molecular weights of several chlorine compounds are 
found by the vapor density method. The compounds are 
analyzed to find the number of grams of chlorine in the 
number of grams of the compound equal to the determined 
molecular weight. And the highest common factor of 
these weights of chlorine is taken as the atomic weight of 
the element. A concise view of the method is shown in 
the following 





H. C. F. 

Hydrochloric acid .... 
Chlorine peroxide .... 
Cyanogen chloride .... 
Chlorine sas 

6l. S 



i x 35-5 
i x 35-5 
i x 35.5 

2 X 3C C 

Chlorine monoxide . . . 
Phosphorus trichloride . . 

I IQ.1 


* JJ'J 
2 X 35.5 

3 x 35.5 
? x "K.c( 

Carbon tetrachloride . . . 



4 x 35.5 

Thirty-five and five tenths is therefore selected as the 
approximate atomic weight of chlorine. 

Atomic weights can also be determined by analysis if 
we know the proportion in which the atoms combine to 
form a molecule of the compound analyzed. Thus, the 
Belgian chemist, Stas, who made masterly determinations 
of atomic weights, found that 121.4993 gm. of silver 

172 Descriptive Chemistry. 

chloride were formed by burning 91.462 gm. of silver 
in chlorine. He knew that one atom of silver and one 
of chlorine unite to form silver chloride ; he also accepted 
35.453 as the atomic weight of chlorine. Hence, he calcu- 
lated the atomic weight of silver thus 

121.4993-91.462= 30.0373, 

which is the weight of the chlorine used. 

91.462 : 30.0373 : : x : 35-453, * = 107.95, 

the atomic weight of silver. 

Approximate atomic weights of the solid elements, espe- 
cially the metals, are checked by applying the law of spe- 
cific heats. This law was announced by Dulong and 
Petit in 1819. It is stated as follows :- 

The product of the specific heat and atomic weigJit of tJie 
solid elements is a constant quantity. 

By Specific heat we mean the quantity of heat necessary 
to raise the temperature of a substance one degree com- 
pared with the quantity necessary to raise the temperature 
of the same weight of water one degree. If the same 
quantity of heat is imparted to equal weights of water and 
mercury, the temperature of the mercury will be much 
higher about 32 times higher than that of the water. 
That is, the mercury requires only about ^ as much heat 
as the water. In other words, the specific heat of mercury 
is ^2> or o -3 r 2 - The specific heat of other elements is simi- 
larly found. 

The constant quantity found by multiplying the specific 
heat by atomic weight is approximately 6.25. This rela- 
tion is illustrated by the following 

Determination of Atomic Weights. 173 







O I7O 


6 8 



6^ 6 

6 04. 


O.I 14. 



Potassium . 
Sodium .... 







6 7^ 









6 cj. 


O OQ4. 

6c A 

6 i c 

' w y i f 


The use of this law in checking atomic weights may be 
illustrated as follows : The specific heat of silver is found 
by experiment to be 0.057; if 6.25 is divided by this num- 
ber, the quotient is approximately 109. This result agrees 
approximately with 108 the accepted atomic weight of 
silver. Again, the specific heat of mercury is 0.0312; if 
6.25 is divided by this number,- the quotient, 200, indicates 
that the atomic weight of mercury is 200 a value obtained 
by other methods. This law has been of assistance in the 
final selection of the approximate atomic weight of several 
elements. Thus, the atomic weight of uranium was finally 
accepted as about 238 instead of 119. Both values agreed 
with analyses, but only the former conformed to Dulong 
and Petit's law. 

The plan followed in determining the atomic weight of zinc illustrates 
the methods actually used. 

(a) When zinc interacts with dilute hydrochloric or sulphuric acid, 
hydrogen is liberated ; and if a known weight of zinc is used, the weight 
of zinc needed to liberate I gm. of hydrogen is easily calculated. 
This number, as we have already seen, is the equivalent of zinc (see 
Equivalents, Chapter IX). Now if one atom of zinc replaces one atom 

174 Descriptive Chemistry. 

of hydrogen, then the atomic weight of zinc and the atomic weight of 
hydrogen will have the same ratio as the weight of zinc and the weight 
of hydrogen found by experiment. According to experiment the 
equivalent of zinc is about 32.5. This is its relation, atom for atom, 
to hydrogen, and, thus far, is its atomic weight. 

() When zinc and hydrochloric acid interact, zinc chloride is 
formed. If it is analyzed, the proportion of zinc to chlorine is about 
32.5 to 35.5. If the elements combine, atom for atom, the atomic 
weight of zinc is 32.5 (assuming that 35.5 is the atomic weight of 

(<:) When zinc is burned in air, zinc oxide is formed. If this com- 
pound is analyzed, the proportion of zinc to oxygen is about 65 to 16. 
If the elements combine atom for atom, the atomic weight of zinc is 
about 65 (assuming that 16 is the atomic weight of oxygen). 

(df) According to these three determinations, the atomic weight of 
zinc is 32.5 or 65. We have assumed that the elements unite atom for 
atom in each compound. This is an incorrect assumption, because an 
atom of zinc cannot have two different weights 32.5 and 65. If the 
atomic weight is 32.5, zinc oxide must consist of one atom of oxygen 
and two of zinc. But if the atomic weight is 65, zinc chloride must 
consist of two atoms of chlorine and one of zinc, and two atoms of 
hydrogen must have been replaced by one of zinc. 

(e) The molecular weight of zinc chloride is found by the vapor 
density method to be about 133. If zinc chloride consists of two 
atoms of chlorine and one of zinc (weighing 65), its molecular weight 
is about 136. In other words, it is evident that our assumption regard- 
ing the number of atoms in zinc chloride is highly probable. 

CO We are not absolutely positive, however, that the zinc in a 
molecule of zinc chloride may not be one atom weighing 65, or two 
atoms weighing 32.5 each. But the atomic weight of zinc determined 
by applying the law of specific heats is 664 (i.e. 6.25 -:- 0.094). This 
shows clearly that the atomic weight of zinc is approximately 65. 

Molecular Formula. In Chapter IX a method was 
given for finding the simplest formula of a compound, viz., 
by dividing the percentage of each element by its atomic 
weight. But the simplest formula is not always the mo- 
lecular formula ; that is, it does not always express the 
composition and number of atoms in a molecule of the 

Molecular Equations. 175 

compound in the gaseous state. Every formula, however, 
is designed to be a molecular formula. Since the molecu- 
lar weight of a compound is twice its vapor density, the 
molecular formula can be calculated from the simplest 
formula. Thus, the simplest formula of a compound of 
carbon and hydrogen was found to be CH 2 . Its vapor 
density was found to be 81.4. Hence its molecular weight 
must be 162.8, which is nearly twelve times that corre- 
sponding to CH 2 . Therefore the molecular formula is 
C 12 H 24 . Molecular formulas of other compounds may be 
similarly found. 

Molecular Equations. Equations which represent re- 
actions between gases are sometimes written as molecular 
equations. Such equations represent changes as taking 
place between the smallest possible physical units, that is, 
between molecules. The molecular equation for the for- 
mation of water from hydrogen and oxygen is 

2 H 2 + O 2 = 2 H 2 O. 

It is read thus : Two molecules of hydrogen unite with one 
molecule of oxygen to form two molecules of water. Since 
most elementary gases consist of molecules, such an equa- 
tion is strictly correct. It should be noted, however, that 
the proportions are the same as in the simpler form of the 
equation. For practical purposes the molecular equation 
is preferable only in the case of gases. 

Molecular equations are sometimes called volume or gas equations, 
because such equations tell at a glance the volumes involved in the re- 
action. Thus- H2 + Cl2 = 2HC1 

means that one volume each of hydrogen and chlorine unite to form two 
volumes of hydrochloric acid gas. This equation is sometimes writ- 

ten H 2 + C1 2 = 2HC1 

I VOL I VOl. 2 VOl. 

176 Descriptive Chemistry. 

Valence. An examination of many formulas obtained 
by the principles just discussed shows certain regularities. 
Take, for example, some binary compounds of hydrogen. 
They fall into four groups, thus 

I. II. III. IV. 

HC1 H 2 O H 3 N H 4 C 

HBr H 2 S H 3 P H 4 Si 

Obviously, the atoms of these elements differ in their 
power of combining with hydrogen atoms. Some unite 
with one atom, some with two atoms, and so on. Atoms 
of other elements besides those in the above list differ in 
their combining power. The power of atoms of an ele- 
ment to hold in combination a certain number of other 
atoms is called the valence or quantivalence of the 
element. The valence of hydrogen is always one. Ele- 
ments which combine atom for atom with one atom of 
hydrogen have the valence one, and are called univalent 
elements or monads ; sodium and potassium are always 
univalent, and so is chlorine in hydrochloric acid. Ele- 
me'nts which combine with two atoms of hydrogen have 
the valence two, and are called bivalent elements or 
dyads ; oxygen, magnesium, and sulphur are bivalent 
elements. So, also, some elements like aluminium, are tri- 
valent or triads ; others, like carbon and silicon, are 
quadrivalent or tetrads; and some, like the nitrogen in 
nitric acid, are quinquivalent or pentads. Elements of 
the same valence combine with or replace each other atom 
for atom. Thus, one atom of sodium replaces one atom 
of hydrogen in hydrochloric acid ; and one atom of oxygen 
combines with one atom of magnesium. Elements of dif- 
ferent valence form compounds in which, as a rule, the 
number of atoms is such that the valences balance, Thus, 

Valence. 177 

a dyad combines with two monads (as in H 2 O), a triad with 
three monads (as in NH 3 ), two triads with three dyads (as 
in A1 2 O 3 ), one tetrad with two dyads (as in CS 2 ), and so 
on. Such compounds, in which the capacity for further 
union has ceased, are said to be saturated or to have no 
free bonds. Compounds in which the valence is not bal- 
anced, or in which free bonds exist, are called unsaturated 
(see Ethylene). 

The valence of an element is always the same in the 
same compound, but it often varies. Thus, the valence of ni- 
trogen is one in N 2 O, two in NO, three in N 2 O 3 , four in NO 2 , 
and five in HNO 3 . Hydrogen, as stated above, always 
has a valence of one ; it is also believed that the valence of 
oxygen is always two. If an element forms no hydrogen 
compound, its valence is determined from compounds con- 
taining elements which are univalent, such as chlorine, 
bromine, and sodium. 

The valence of elements in saturated compounds of two 
elements is easily deduced from the formula, because in 
such compounds the total valence of all the atoms of each 
element must be equivalent Thus in the formula CaO, 
the valence of calcium is two, because the single atom of 
calcium is combined with a single atom of a bivalent 
element. The valence of phosphorus in P 2 O 5 is five, be- 
cause the two atoms furnish a total valence of ten, which 
is required by the five atoms of the bivalent element 
oxygen. In CH 4 the valence of carbon is four, because the 
single atom is combined with four atoms of hydrogen. 

Radicals have a valence, since in chemical changes they 
act like atoms. The valence of ammonium (NH 4 ) is one, 
and of hydroxyl (OH ) is one. Thus, NH 4 C1 is the formula 
of ammonium chloride, NaOH of sodium hydroxide, but 
Ca(OH) 3 of calcium hydroxide. 


1 7 8 Descriptive Chemistry. 

The valence of elements in unsaturated compounds can- 
not be told by mere inspection ; a knowledge of the prop- 
erties of the compound is necessary. So also the valence 
of some elements in compounds containing three or more 
elements is not readily told from the formulas^ some 
knowledge of the methods of formation, relations to other 
compounds, and general properties is needed. A discus- 
sion of these principles is beyond the scope of this book. 
However, in the case of most acids, bases, and salts, an 
arbitrary rule may be cited. In these compounds the total 
valence of the oxygen atoms balances the total valence of 
the other elements. Thus, in nitric acid, HNO 3 , the va- 
lence of nitrogen is nve, while in nitrous acid, HNO 2 , it is 

Some chemists prefer to regard valence as the quotient obtained by 
dividing the atomic weight by the equivalent weight. For example, 
the valence of oxygen is 2 the quotient of 16 -4- 8. Such a view is 
not inconsistent with the one generally held, because valence is the 
direct outcome of. composition. 

The valence of elements may be represented in several ways, e.g. 

H', H , O , O = , N . Sometimes formulas are written to 

show the valence, e.g. / H 

Hydrochloric acid, H - Cl, Water, H - O - H, Ammonia, N - H. 

Such formulas are called structural or graphic formulas to distinguish 
them from the ordinary or empirical formulas. Structural formulas 
are not intended to show how the atoms are arranged in space. We 
know very little about the space relations of atoms. They simply indi- 
cate certain relations not shown by the empirical formulas. They are 
especially helpful in organic chemistry (see Chapter XXXI). 


i . Review (a) Boyle's law, and (ft) Charles's law. 

2. State and illustrate Gay-Lussac's law. 

3. Give a brief account of (a) Gay-Lussac ? () Avogadro, (c) Stas, 



4. State and illustrate Avogadro's hypothesis. 

5. What is the relation of the molecular weight of a gas to (a) the 
molecular and () the atomic weight of hydrogen ? 

6. (a) State the argument proving that a molecule of hydrogen con- 
sists of two atoms, (b} Apply the same argument to oxygen. 

7. What is the relation between molecular weight and vapor den- 
sity ? Illustrate your answer. What application is made of this 
relation ? 

8. Why is the formula of water H 2 O and not HO or H 2 O 2 ? 

9. Why is the formula of ozone O 3 ? 

10. (a) How are molecular weights determined ? () How are 
atomic weights found from molecular weights ? 

11. Illustrate the method of determining atomic weights by chemical 

12. What is a molecular formula ? What is the molecular formula 
of oxygen, nitrogen, chlorine, and hydrogen ? How is a molecular 
formula determined ? Illustrate your answer. 

13. What is a molecular equation ? Give two illustrations. How 
does it differ from an ordinary chemical equation ? Of what use are 
such equations ? 

14. Define () valence, (b} monad, dyad, triad, tetrad, pentad, 
(c) univalent element, bivalent element, (d) saturated compound, 
(e) unsaturated compound. 

15. What is the valence of hydrogen ? Why ? Of oxygen ? Why ? 
How may valence be found by inspecting a binary formula ? What is 
the valence of NH 4 and OH ? 

1 6. Illustrate the ways valence may be represented. 

17. Distinguish between structural and empirical formulas. 

1 8. What is the valence of sodium in (a} sodium chloride, (b) so- 
dium nitrate (NaNO 3 ), (c} sodium sulphate (Na 2 SO 4 ), (d) sodium 
hydroxide (NaOH) ? 

19. What is the valence of sulphur in (a) snlphur dioxide (SO 2 ), 
(b) sulphur trioxide (SO 3 ), (c} hydrogen sulphide (H 2 S), (d) sulphuric 
acid, (e) copper sulphate (CuSO 4 ) ? (Suggestion. In oxygen acids, 
the oxygen valence balances the sum of the valence of the other 

20. What is the valence of (#) aluminium in aluminium oxide 
(Al 2 Oo), (b} carbon in carbon tetrachloride (CC1 4 ), (c) phosphorus in 
phosphorus pentoxide (P 2 O.) ? 


180 Descriptive Chemistry. 

21. What is the valence of (a} silver and chlorine in silver chloride 
(AgCl), () calcium and chlorine in calcium chloride (CaCl 2 ), (<;) oxy- 
gen in water, (d) oxygen and calcium in calcium oxide or lime (CaO) ? 


1. The vapor densities of certain gases is as follows : (#) hydro- 
chloric acid 18.25, (b} chlorine 35.5, (c) ammonia 8.5, (d) nitrogen 14, 
0) steam 9. Calculate the molecular weight of each. 

2. Calculate the simplest formula of the compounds which have 
the indicated composition: (a} N = 82.353, H = 17.647; () O = 30, 
Fe (iron) =70; (c) H = i, C = 12, K (potassium) =39, O = 48. 

3. A liter of sulphurous oxide gas (SO 2 ) weighs 2.8672 gm. 
What is the molecular weight of this compound ? 

4. If 1500 cc. of carbon monoxide gas (CO) weigh 1.8816 gm., 
what is the molecular weight of the compound ? 

5. Calculate the molecular formula of the compounds corresponding 
to the following data: (a) C = 73.8, H = 8.7, N = 17.1, vapor density 
= 80.2; () C=92.3, H = 7.7, vapor density =38.8 ; (c) C = 39.9, 
H = 6.7, O =53.4, vapor density = 30.5. 

6. What volumes of factors and products are represented by the 
equations (a) H 2 + C1 2 = 2 HC1, () 2 H, + O 2 = 2 H,O, (c) 3 H, + 
N 2 = 2 NH 3 , (d) N 2 + O 2 = 2 NO, (e) 2 NO + O, = 2 NO, ? 

7. If 20 1. of hydrogen are allowed to interact with 10 1. of chlo- 
rine, (a) how many liters of hydrochloric acid gas are produced, and 
() which gas and how much remains ? 

8. How many liters of hydrogen gas can be obtained from 4 1. 
of hydrochloric acid gas ? 

9. If 91.462 gm. of silver, when heated in chlorine, yield 121.4993 gm. 
of silver chloride, what is the atomic weight of chlorine ? (Assume 
Ag = 108.) 

10. How many liters of the component gases can be obtained by 
the decomposition of 6 1. of ammonia gas ? 

11. Find the simplest formulas of the substances having the follow- 
ing composition : (a) H = 1.58, N = 22.22, O = 76.19 ; (^) O = 47.52, 
N = 13.86, K = 38.61. 

12. A certain weight of copper oxide, when heated in a current 
of hydrogen, lost 59.789 gm. of oxygen and formed 67.282 gm. of 
water. (a) If O = 16, what is the atomic weight of hydrogen ? 
() If H = i, what is the atomic weight of oxygen ? 


Occurrence of Carbon. Uncombined carbon is found 
pure in nature as diamond and graphite ; in a more or less 
impure state it occurs as coal and similar substances, 
which are included in the term amorphous carbon. Car- 
bon forms a vast number of compounds, natural and 
artificial. Combined with hydrogen and oxygen, and 
occasionally with nitrogen also, it is an essential constitu- 
ent of plants and animals. Meat, starch, fat, sugar, wood, 
cotton, paper, soap, wool, wax, flour, albumen, and bone 
contain carbon. It is also a component of carbon dioxide 
and of carbonates, such as limestone, chalk, and marble. 
Illuminating gases, kerosene and other products of petro- 
leum, turpentine, alcohol, chloroform, ether, and similar 
liquids are compounds of carbon. It is estimated that 
0.22 per cent of the weight of the earth's crust is carbon. 

Diamond is pure crystallized carbon. It is found in 
only a few places in the earth. When taken from the 
mine, diamonds are rough-looking stones ; some are crystals, 
some are rounded like peas, and many are irregular ; they 
must be cut and polished to bring out the luster and make 
them sparkle (Fig. 24). The highly prized diamonds are 
colorless and without a flaw, and are said to be "of the 
first water " ; yellow ones from South Africa are common, 
and occasionally a blue, pink, red, or green one is found ; 
a very impure variety is black. 



Descriptive Chemistry. 

The diamond is insoluble in all liquids at the ordinary 
temperature, has the high specific gravity of 3.5, and is the 
hardest known substance. 

It is brittle and may be shattered by a blow with a 


FIG. 24. Diamonds. 


Diamonds have always been prized as gems on account of their 
beauty, rarity, and permanency. Besides being worn as jewels, they 
are used to cut glass, and the powder and splinters (known as bort) 
are used to grind and polish diamonds and other hard gems. The im- 
pure variety which comes from Brazil, and is called carbonado, is set into 
the end of the " diamond drill,' 1 which is used extensively for boring 
artesian wells and drilling hard rocks. 

The diamond was formerly found in gravel deposits in India, and in 
later years in Brazil. Since 1867, however, about 95 per cent of the dia- 
monds of commerce have come from South Africa. They occur in a 
bluish volcanic rock along the Vaal River, and especially near Kimberley. 
Over eight tons of diamonds have been found in South Africa in the 
last twenty-five years ! 

The successive investigations of Lavoisier, Dumas, and Davy, ex- 
tending from 1772 to 1814, showed that diamond is carbon, for when 
pure diamond was burned in oxygen, the only product was carbon 

dioxide. This result, which ad- 
mits of no doubt, has been verified 
by many famous investigators. 
Diamonds have been made by 
Moissan. He dissolved pure char- 
coal in melted iron, and poured the 
molten mass into water. The sur- 
face was so suddenly cooled that a tremendous pressure was exerted 

FlG. 25. Artificial diamonds (enlarged) 
prepared by Moissan. 

Carbon and its Oxides. 183 

by the expanding iron inside the crust. This pressure caused the cool- 
ing carbon to crystallize into diamond. The crystals were very small, 
most of them were black, a few were white, but all had the properties 
of the diamond (Fig. 25). 

Large diamonds have a fascinating history, since most of them have 
passed through many hands before finding a place among royal jewels. 
The largest is the Orloff, which weighs 194! carats, and is in the scepter 
of the Czar of Russia. 1 The Kohinoor, which now weighs about 106 
carats, is one of the crown jewels of England. 

Graphite is a soft, black, shiny solid, which is smooth 
and soapy to the touch. Pure graphite is carbon. It occurs 
native in large quantities and in many places. One va- 
riety is found in abundance at Ticonderoga, New York. 
Other famous localities are Ceylon, eastern Siberia, Bava- 
ria, and Italy. Sometimes crystals and grains are found, 
but it usually occurs in flaky masses or slabs. Unlike 
diamond, graphite is a good conductor of electricity and is 
often used to coat moulds in electrotyping. It is so soft 
that it blackens the fingers and leaves a black mark on 
paper when drawn across it. This property is indicated 
by the name graphite, which is derived from a Greek word 
(grap/iein) meaning to write. It resembles diamond in its 
insolubility in liquids at the ordinary temperature. Its 
specific gravity is 2.2, being considerably lighter than dia- 
mond. It produces only carbon dioxide when burned in 
oxygen ; but unlike diamond, it turns into carbon dioxide 
by heating to a very high temperature in the air. Graphite 
was once supposed to contain lead, and rs even now often 
incorrectly called " black lead " and plumbago. It is used 
to make stove polish and protective paints, as a lubricant 
where oil cannot be used, as' the principal ingredient of 

1 A carat equals 3J Troy grains (or 0.205 gm.). The term is derived from the 
carob bean, which was used for ages by the diamond merchants of India as a 
small weight. 


Descriptive Chemistry. 

graphite crucibles, in which metals are often melted, and 
in making electrodes for the huge electric furnaces. 

Immense quantities of graphite are consumed in the manufacture of 
lead pencils. The graphite is washed free from impurities, ground to a 
fine powder, mixed with more or less clay, and then pressed through 
perforated plates, from which the "lead" issues in tiny rods. These 
are dried, cut into the proper lengths, baked to remove all traces of 
moisture, and then inserted in the wooden case. 

In the United States in 1902 over four million pounds of graphite 
were mined, and over thirty-two million pounds were imported. 

Molten iron and other metals dissolve carbon, and when the metals 
cool the carbon crystallizes as graphite. Moissan incidentally obtained 
considerable graphite in making diamonds. Artificial graphite is now 
a commercial article (see Chapter X). 

Amorphous Carbon is a broad term, including all vari- 
eties of coal and charcoal, lampblack, and gas carbon. 
They are the non-crystalline forms of impure carbon. 
The word amorphous means literally "without form," and 
it is often used to designate soft, powdery, and uncrys- 
tallized substances. 

Coal is a term applied to several varieties of impure carbon. It may 
be regarded as the final product derived from vegetable matter by heat 
and pressure to which it was subjected through long geological periods. 

Ages ago the vegetation was exceedingly dense and luxuriant upon 
land slightly raised above the sea. In process of time this vegeta- 

FiG. 26. Section of part of the earth's crust near Mauch Chunk, Penn., 
showing layers of coal. 

tion decayed, accumulated, and slowly became covered with sand, mud, 
and water. The heat of the earth and the enormous pressure of the 
overlaying deposits changed the vegetable matter into more or .less 

Carbon and its Oxides. 185 

impure carbon. This series of geological and chemical changes was 
repeated, and as a result we find in the earth layers or seams of carbo- 
naceous matter varying in thickness and composition (Fig. 26). These 
are the coal beds. 

Coal .beds contain proofs of their vegetable origin, viz., impressions 
of vines, stems, and leaves of plants, and similar vegetable substances 

FIG. 27. Fossil found in a FIG. 28. Section of coal as seen through 

coal bed. a microscope. 

(Fig. 27). A thin section of coal examined through a microscope re- 
veals a distinct vegetable structure (Fig. 28). 

There are three principal kinds of coal, (i) Bitumi- 
nous or soft coal is used to make illuminating gas, coke, 
and as a fuel for steam ; it burns with a smoky flame, and 
in burning produces much volatile matter. (2) Anthra- 
cite coal is hard and lustrous. It ignites with difficulty, 
burns with little or no flame, and produces an intense heat. 
It is used mainly for domestic purposes, heating and 
cooking, especially in eastern United States. (3) Lig- 
nite or brown coal is the least valuable as fuel. It often 
shows the woody fiber and was probably formed much 
later than the other varieties. Peat, strictly speaking, is 
not coal, though it is used as fuel in some places, espe- 
cially in Ireland and Holland. It is formed by the slow 


Descriptive Chemistry. 

decay of roots and other vegetable matter under water, 
and represents an early stage of coal formation. 

The average composition of different kinds of coal is 
seen by the following table : 

, KIND. 





Lignite .... 


2O Q 

jO 2 


-> w -y 

Bituminous .... 




Anthracite .... 




Some anthracite coals contain as much as 95 to 99 per 
cent of carbon, and some bituminous coals as little as 65 
per cent. Peat and wood contain still less carbon, but 

FIG. 29. Coal fields in the United States. 

more volatile matter. The volatile matter includes nitro- 
gen, hydrogen, and sulphur. These facts show that vege- 
table matter, in passing through the changes which finally 

Carbon and its Oxides. 

end in coal, loses volatile matter, Anthracite coal, which 
is found at different depths and associated with rocks of 
different ages, shows that it was formed from the bitumi- 
nous variety by the great pressure caused by mountain 
building. Hence it loses volatile matter and becomes hard. 

Coal is widely distributed in the crust of the earth, but the deposits 
vary in extent and quality. It underlies about one sixth of the area of 
the United States, the anthracite variety covering less than five hundred 
square miles in eastern Pennsylvania (Fig. 29). The United States 
now leads the world in coal production, 
furnishing about one third of the total 
supply. England for many years headed 
the list, and even now furnishes a large 
amount, for its deposits are extensive 
(Fig. 30). 

Charcoal is a variety of amor- 
phous carbon obtained by heating 
wood, bones, ivory, and other 
organic matter in closed vessels, 
or by partially burning them in 
the air. Th? process consists 
essentially in driving off the vola- 
tile matter and retaining the 

Wood Charcoal is a black, 
brittle solid, and often has the 
form of the wood from which it 
is made. It is insoluble, though 
its mineral impurities may be removed by acids. It 
burns without lame or much smoke, and leaves a white 
ash. The compact varieties conduct heat and electricity, 
but porous charcoal is a poor conductor. It resists the 
action of many chemicals; hence fence posts, telegraph 
poles, and wooden piles are often charred before being 


FlG. 30. Coal deposits in the 
British Isles. 

1 88 Descriptive Chemistry. 

put into the ground. Most varieties are very porous, and 
when thrown upon water charcoal floats, owing to the 
presence of air in its pores. Its porosity makes charcoal 
an excellent absorber of gases, some varieties absorbing 
ninety times their bulk of ammonia gas. Sewers and foul 
places are sometimes purified by charcoal. It will also 
absorb colored substances from solutions. This is espe- 
cially true of animal charcoal (see below). Foul air and 
water may be partially purified by charcoal, which forms 
the essential part of many water filters in houses. Char- 
coal used for such a purpose, however, must be renewed 
or often heated to redness; otherwise it becomes clogged 
and contaminated. Charcoal is never pure carbon, the 
degree of purity depending upon the kind of wood used, 
as well as the temperature and method employed. 

Besides the uses of charcoal mentioned above, it is used 
as a fuel, in the manufacture of steel and of gunpowder, 
and as a medicine. It reduces oxides when heated with 
them, thus 

2 CuO + C = 2 Cu + CO 2 
Copper Oxide Carbon Copper Carbon Dioxide 

Wood charcoal is made either in a charcoal pit or kiln, or in a large 
retort. Where wood is plentiful, it is loosely piled into the shape 
shown in Figure 31, and covered with turf to prevent free access of air, 
though small holes are left at the bottom and a larger one at the top of 
a central flue, so that sufficient air can pass through the pile. The 
wood is lighted, and as it slowly burns care is taken to regulate the 
supply of air, so that the wood will smolder but not burn up. The 
volatile matter escapes and charcoal remains, the average yield being 
about 20 per cent of the weight of the wood. This method is crude, 
uncertain, and wasteful. Much charcoal is now made by heating 
wood in closed retorts, no air whatever being admitted. By this 
method, which is called dry or destructive distillation, the yield of 
charcoal is 30 per cent and all the volatile matter is saved. In the 

Carbon and its Oxides. 189 

ordinary combustion of wood, the hydrogen forms water and the oxy- 
gen forms carbon dioxide ; but in dry distillation, where no oxygen is 
present, much of the hydrogen forms volatile compounds with the car- 
bon and oxygen. Among these volatile products are methyl alcohol 

FIG. 31. Wood arranged for burning into charcoal. 

and acetic acid. These are commercial substances, and contribute to 
the profit of the process. More or less charcoal is obtained by heating 
any compound of carbon, e.g. sugar or starch, the charring being a test 
for carbon. 

Animal Charcoal or Bone Black is made by heating bones in a closed 
vessel, and by heating a mixture of blood and sodium carbonate. It 
contains only about 10 per cent of carbon, but this carbon is dis- 
tributed throughout the porous mineral matter of the bone, which is 
almost entirely calcium phosphate. Under the name of ivory black, 
animal charcoal is used as a pigment, especially in making shoe-black- 
ing. It is extensively used to remove the color from sugar sirups, oils, 
and other liquids colored by organic matter. 

Coke is made by expelling the volatile matter from soft 
coal, somewhat as charcoal is made from wood. It is left 
in the retorts when coal is distilled in the manufacture of 
illuminating gas. On a large scale it is made by heating 
a special grade of soft coal in huge brick ovens, shaped 
like a beehive, from which air is excluded after combus- 
tion begins. Sometimes the coke is made in closed retorts 
constructed so as to save the by-products, ammonia, tar, 

190 Descriptive Chemistry. 

organic compounds, and combustible gases. This method 
not only yields more coke, but is also more profitable be- 
cause the by-products are sold and the combustible gas is 
used to heat the retorts. Coke is a grayish, porous solid, 
harder and heavier than charcoal. It burns with no smoke 
and a feeble flame. It contains about 90 per cent of car- 
bon, the rest being the mineral matter originally in the coal. 

Immense quantities of coke are used in the manufacture of iron and 
steel. It is superior to coal for this purpose, because it gives a greater 
heat when burned, reduce's oxides easily, and contains little or no 
sulphur or other substances harmful in the iron industries. Coke is 
the fuel used in making nine tenths of the pig iron in the United 
States, and over twelve million tons (or about three fourths of the 
total amount) are made annually in the Connellsville district, near 
Pittsburg, Pennsylvania. 

Gas Carbon is amorphous carbon which is gradually deposited upon 
the inside of the retorts used in the manufacture of illuminating gas. 
It is a black, heavy, hard solid, and is almost pure carbon. It is a good 
conductor of electricity, and is extensively used for the manufacture of 
the carbon rods of electric lights and for plates of electric batteries. 

Lampblack is prepared by burning oil or oily substances rich in 
carbon in a limited supply of air. The dense smoke, which is mainly 
finely divided carbon, is passed through a series of condensing cham- 
bers, where it is collected upon coarse cloth or a cold surface. Its 
formation is illustrated on a small scale by a smoking lamp, and the 
soot deposited is the same as lampblack. Lampblack is one of the 
purest forms of amorphous carbon, and it is used in making printer's 
ink and certain black paints. 

Allotropism. Diamond, graphite, and amorphous car- 
bon, though exhibiting essentially different properties, are 
identical in composition. All are carbon. They can be 
changed into one another, the amorphous form into graph- 
ite and finally into diamond and the diamond into amor- 
phous carbon. Each burns in oxygen and the product is 
carbon dioxide. Furthermore, the same weight of each 

Carbon and its Oxides. 191 

forms the same weight of carbon dioxide, i.e. when 12 
gm. of each are burned, 44 gm. of carbon dioxide are 
always produced. There is no doubt about their identity, 
though no one has explained it. The property of assum- 
ing more than one elementary form is called allotropism 
or allotropy (from Greek words meaning another form). 
The more uncommon form is called an allotrope or an 
allotropic modification of the other. It is believed by some 
that allotropism is due to a difference in the number of 
atoms in a molecule of the element. 


Carbon and Oxygen do not unite at the ordinary tem- 
perature. But when carbon is heated in air, in oxygen, or 
with some oxides, carbon dioxide (CO 2 ) is formed ; if the 
supply of oxygen is limited, then carbon monoxide (CO) 
is formed. 

Occurrence and Formation of Carbon Dioxide. The 

occurrence of carbon dioxide in the atmosphere and in 
many natural waters has already been mentioned. It is 
the main product of ordinary combustion, respiration of 
animals, and decay. In all these processes the carbon 
comes from organic matter, while the oxygen comes from 
the air, from the organic matter, or from both. 

Ordinary combustion is a chemical combining of carbon 
and oxygen. Hence, when carbon or a substance contain^ 
ing it is burned, carbon dioxide is formed. The equation 
for this change is 

C + 2 C0 2 

Carbon Oxygen Carbon Dioxide 

Carbon dioxide is formed by the combustion of such com- 
mon substances as wood, coal, charcoal, coke, oils, waxes, 

192 Descriptive Chemistry. 

cotton, bone, starch, sugar, meat, bread, alcohol, camphor, 
and illuminating gas. 

The continuous oxidation of the tissues and foods in the 
body produces carbon dioxide (see Relation of Oxygen to 
Life). And if we exhale the breath through a glass tube 
into limewater, the carbon dioxide which is in the breath 
turns the limewater milky the usual test for carbon 
dioxide. The equation for the change is 

CO 2 + Ca(OH) 2 = CaCO 3 + H 2 O 

Carbon Dioxide Limewater Calcium Carbonate 

When vegetable and animal matter decays, carbon 
dioxide is formed. Many kinds of organic matter fer- 
ment, especially those containing sugar. By alcoholic 
fermentation the sugar changes into carbon dioxide and 
alcohol (see Alcohol), thus 

C 6 H 126 = 2CO 2 + 2C 2 H 6 O 

Sugar Carbon Dioxide Alcohol 

The Preparation of Carbon Dioxide is usually accom- 
plished by the interaction of a carbonate and an acid. 
Calcium carbonate (limestone or marble) and hydrochloric 
acid are usually used. The operation may be easily per- 
formed in any glass vessel by pouring the acid upon the 
carbonate. The equation for the chemical change is 

CaCOg + 2HC1 = CO 2 + CaCl 2 + H 2 O 
Calcium Carbon Calcium 

Carbonate Dioxide Chloride 

This gas may also be prepared by heating matter con- 
taining carbon, or by strongly heating carbonates (as in 
making lime), thus 

CaCOg CO 2 + CaO 

Calcium Carbonate Carbon Dioxide Lime 

Carbon and its Oxides. 

Properties of Carbon Dioxide. This gas has many 
important properties besides those mentioned under The 
Atmosphere. It has a slight taste and odor, but no color. 
It is one and a half times heavier than air, and a liter 
under standard conditions weighs 1.977 gm. On ac- 
count of its weight it can be collected by downward dis- 
placement and poured from one vessel to another. For 
the same reason, it is often found at the bottom of old or 
deep wells, in some valleys near lime kilns or volcanoes, 
and in mines after explosions. At the ordinary tempera- 
ture and pressure, water dissolves its own volume of 
carbon dioxide. Under increased pressure more gas dis- 
solves, which escapes readily when the pressure is re- 
moved. Hence " soda water," which is made by forcing 
carbon dioxide into water, effervesces and froths when 
drawn from the soda fountain. Many natural waters and 
manufactured beverages (such as champagne and beer) 
sparkle and effervesce for the same reason. This gas may 
be liquefied by subjecting it to high pressure and low 
temperature. It was first liquefied by Faraday by the 
method used for chlorine. Liquid carbon dioxide is now 
made in large quantities by forcing the gas into steel 
cylinders by powerful pumps, the gas being obtained in 
many cases from the fermenting vats of breweries. 
When a cylinder of liquid .carbon dioxide is opened, the 
liquid evaporates so rapidly that a portion of it becomes 
a white, snowlike solid. Both the liquid and solid carbon 
dioxide are articles of commerce, and are sometimes 
used to prepare "soda water," to extinguish fires, to 
improve wines, and to produce very low temperatures. 
Carbon dioxide extinguishes burning objects, such as a 
blazing stick or lighted candle; indeed, air containing 
from 2.5 to 4 per cent of carbon dioxide will extinguish 

194 Descriptive Chemistry. 

small flames. Hence the gas is often used to extinguish 
fires. Many small fire extinguishers contain sodium 
carbonate and sulphuric acid, so arranged that when 
desired, carbon dioxide gas may be generated from them 
under pressure. A stream of the gas forced upon a 
small blaze will often prevent a serious fire. In other 
forms, the carbon dioxide, which is similarly generated, 
forces water from the extinguisher. 

Relation of Carbon Dioxide to Life. Animals die when 
put into carbon dioxide. It cuts off the supply of oxygen 
as water does from a drowning man. The presence of a 
small quantity in the air is objectionable, since it is said to 
produce headache and drowsiness; but much of the dis- 
comfort felt in badly ventilated rooms and attributed to 
carbon dioxide is doubtless due to water vapor, and to 
poisonous substances produced from the organic mat- 
ter exhaled from the lungs. On the other hand, carbon 
dioxide is an essential food of plants. Through their 
leaves and other green parts they absorb carbon dioxide 
from the atmosphere, decompose it, reject the oxygen, and 
store up the carbon in the form of starch. The sunlight 
and the green coloring matter aid the plant in manufac- 
turing its food out of the water (obtained through the roots 
from the soil) and the carbon of the carbon dioxide ob- 
tained from air. Plants thus serve to keep the atmosphere 
free from an excess of carbon dioxide, the proportion 
present in the air being very small and practically con- 

Carbonic Acid. Carbon dioxide gas is often called carbonic acid gas, 
or simply carbonic acid. It is believed that carbon dioxide, when passed 
into water, combines with the water and forms a weak, unstable acid, 
which is, strictly speaking, carbonic acid. The equation for this 
change is 

Carbon and its Oxides. 195 

CO 2 + H 2 = H 2 CO 3 
Carbon Dioxide Carbonic Acid 

Such a solution reddens blue litmus and decolorizes pink phenolphthal- 
ein. Carbonic acid has never been obtained free, and is so unstable 
that it easily breaks up by gentle heat into carbon dioxide and water, 

H,CO 3 = CO 2 + H 2 0. 

Carbon dioxide is sometimes called carbonic anhydride, to denote its 
relation to the acid. 

Carbonates are salts corresponding to the unstable 
carbonic acid. They are stable compounds. The most 
abundant natural carbonates are those of calcium, magne- 
sium, and iron. Immense quantities of sodium and potas- 
sium carbonates are manufactured. 

A few carbonates are formed by direct combination of an oxide and 
carbon dioxide, but most of them are formed by passing carbon dioxide 
into the corresponding hydroxide, thus 

CO 2 + Ca(OH) 2 CaCO 3 + H 2 O 

Calcium Hydroxide Calcium Carbonate 

Many carbonates are insoluble in water, e.g. calcium carbonate, the 
test for carbon dioxide depending upon this fact. Others, e.g. sodium 
and potassium carbonate, are very soluble. There are two classes of 
carbonates, the normal and the acid. Normal sodium carbonate is 
Na 2 CO 3 , and acid sodium carbonate is HNaCO 3 . The latter is often 
called sodium bicarbonate. Normal calcium carbonate is CaCO 3 , and 
acid calcium carbonate is H 2 Ca(CO 3 ) 2 ; 4he latter is unstable, and is 
easily decomposed by heat into normal calcium carbonate. 

Composition of Carbon Dioxide. If a known weight of pure car- 
bon, such as diamond or graphite, is burned in oxygen, it is found that 
for 12 parts of carbon used there are 44 parts of carbon dioxide formed. 
Hence 12 parts of carbon unite with 32 parts of oxygen. The vapor 
density of the gas is 22, and the molecular weight must be 44. These 
facts necessitate the formula CO 2 . 

196 Descriptive Chemistry. 

History of Carbon Dioxide. This gas was described in the seven- 
teenth century by Van Helmont, who called it gas sylvestre. He 
prepared it by the interaction of acids and carbonates, detected it in 
mineral water, and observed its formation during combustion and fer- 
mentation, as well as its action on animals and flames. Black, in 1755, 
showed that carbon dioxide is essentially different from ordinary air and 
that the gas is readily obtained from magnesium and calcium carbonates. 
Since the gas was combined or " fixed " in these substances, he called 
the gas fixed air. His work was verified in 1774 by Bergman, who 
called the gas acid of air. Lavoisier first proved it to be an oxide of 

Carbon Monoxide is formed when carbon is burned in a 
limited supply of air, thus 

C + O CO 

Carbon Oxygen Carbon Monoxide 

If carbon dioxide is passed over heated charcoal, the prod- 
uct is carbon monoxide. That is, carbon reduces carbon 
dioxide to carbon monoxide, the equation for the change 

C0 2 + C 2 CO 

Carbon Monoxide 

This chemical change takes place in every coal fire. The 
oxygen of the air entering the bottom of the fire unites with 
the carbon to form carbon dioxide ; the latter gas in passing 
through the hot carbon of the fire is reduced to carbon 
monoxide. Some of the carbon monoxide escapes and 
some burns with a flickering bluish flame on the top of 
the fire. 

If steam is passed over red-hot coke or charcoal, a mixture of carbon 
monoxide and hydrogen is produced. This mixture enriched by vapor 
from oils is known as water gas (see Water Gas) . 

Carbon monoxide is usually prepared by gently heating 
a mixture of oxalic acid and sulphuric acid in a flask, and 

Carbon and its Oxides. 197 

collecting the gaseous product over water. The oxalic acid 
decomposes thus 

C 2 H 2 O 4 = CO + CO 2 + H 2 O 

Oxalic Acid Carbon Monoxide Carbon Dioxide 

The carbon dioxide may be removed by passing the mixed 
gases through a solution of sodium hydroxide. 

Carbon monoxide is a gas without color, odor, or taste, and 
is only slightly soluble in water. It burns with a bluish 
flame, forming carbon dioxide, thus 

2 CO + .O 2 2CO 2 

Carbon Monoxide Carbon Dioxide 

Carbon monoxide is extremely poisonous, and it is doubly 
dangerous because its lack of odor prevents its detection in 
time to escape its stupefying effect. Many deaths have 
been caused by breathing air containing it. Carbon mo- 
noxide forms a compound with one of the constituents of 
the blood, and those who have been poisoned by it cannot 
be revived by air, as in the case of suffocation by carbon 
dioxide. It is a constituent of ordinary illuminating gas, 
and care should always be taken to prevent the escape of 
illuminating gas (as well as the gas from a coal stove or 
furnace) into rooms occupied by human beings. At a high 
temperature carbon monoxide unites easily with oxygen, 
and is, therefore, an important agent in the reduction of 
iron ores in the blast furnace. This action might be rep- 
resented thus 

Fe 2 3 + 3 CO = 2Fe + 3 CO 2 

Iron Oxide Carbon Monoxide Iron Carbon Dioxide 

Carbon monoxide, which is sometimes called carbonic oxide, forms no 
acid and therefore no salts. It does not make limewater milky, thus 
being readily distinguished from carbon dioxide. Its blue flame dis- 

198 Descriptive Chemistry. 

tinguishes it from all other gases which burn. It unites directly with 
chlorine to form carbonyl chloride (phosgene, COC1 2 ), and with some 
metals, forming metallic carbonyls, e.g. nickel carbonyl (Ni(CO) 4 ). 

Cyanogen is a compound of carbon and nitrogen having 
the composition corresponding to the formula (CN) 2 . It 
is a colorless gas, has the odor of peach kernels, is exceed- 
ingly poisonous, and burns with a purplish flame. It may 
be prepared by heating mercuric cyanide (Hg(CN) 2 ). 
Cyanogen is a radical, and in compounds it acts like an 
element. Its corresponding acid is hydrocyanic or prus- 
sic acid (HCN). This acid is prepared by heating a 
cyanide with sulphuric acid, "just as hydrochloric acid is 
obtained from a chloride. The solution smells like peach 
kernels, and is one of the most deadly of all known poisons. 
Potassium cyanide is a white, deliquescent solid. It is 
a deadly poison. Large quantities are used in gold and 
silver plating and in the " cyanide process " of extracting 
gold from its ores, as described under that metal. Other 
cyanogen compounds are cyanic acid (CNOH), sulpho- 
cyanic acid (CNSH), and potassium sulphocyanate 
(CNSK). The last is a white, crystallized salt, which 
produces a beautiful red solution when added to certain 
soluble iron compounds, and is therefore used to detect 
this metal. Salts of complex acids related to hydrocyanic 
acid are used in dyeing, many being prepared from the 
most common one potassium ferrocyanide or yellow 
prussiate of potash. They will be described in the chap- 
ter on Iron. 


1. What is the symbol and atomic weight of carbon? 

2. In what forms does free carbon occur in nature? Name ten famil- 
iar solids, three liquids, and two gases containing carbon. What pro- 
portion of the earth's crust is carbon? 

Carbon and its Oxides. 199 

3. What is diamond? How could the correctness of your answer be 
shown? State (a} the source, (b) the properties, and (c) the uses of 
diamonds. Give a brief account of one or more famous diamonds. 

4. What is graphite? What is its chemical relation to diamond, 
and how could this relation be proved ? State (a) the source, (b} the 
properties, and (c) the uses of native graphite. 

5. What is {a} black lead, () plumbago, (c) bort, (d) carbonado, 
(e) native graphite, (/) artificial graphite? 

6. Give a brief account of the manufacture of lead pencils. What 
is the literal meaning of graphite? 

7. Review artificial graphite (see Chapter X). 

8. What does the term amorphous carbon include? Does the car- 
bon in these impure forms differ chemically from diamond and graphite? 

9. How was coal formed? Give several proofs of its origin. State 
the properties and uses of (a) bituminous coal, (b} anthracite coal, and 
(V) lignite. What besides carbon does it contain? Where is coal 
found ? 

10. W T hat is charcoal ? State (a} the properties, and (b} the uses of 
wood charcoal. Give a brief account of both methods of preparing wood 
charcoal. State the preparation, properties, and uses of animal charcoal. 

u. What is coke? How is it made? What are its properties? 
How is it related to the iron industries? 

12. What is gas carbon? What is its source? State its properties 
and uses. 

13. What is lampblack? State its method of preparation, properties, 
and uses. 

14. Define and illustrate (a} amorphous, and (b} allotropism. 

15. Develop the topics: (a} carbon is a reducing agent, (b} carbon 
monoxide is a reducing agent, (c) diamond, graphite, and pure amor- 
phous carbon illustrate allotropism. 

1 6. What is (a) hard coal, (b) soft coal, (c) peat, (d} boneblack, 
(e) soot, (/) lampblack, (g) lignite, (h) electric light carbon? 

17. Give the names and formulas of the two oxides of carbon. How 
is each formed from carbon and oxygen? 

1 8. Describe the occurrence and formation of carbon dioxide. What 
is always obtained by burning a substance containing carbon ? Give 
the simplest equation for this chemical change. 

19. Describe fully the action of carbon dioxide on limewater. Give 
the equation for the reaction. 

2OO Descriptive Chemistry. 

20. What is the relation of carbon dioxide to (a} respiration, (b) fer- 
mentation of sugar, (c) decay, (d) making lime ? 

21. What is the test for (a) carbon, () carbon monoxide, (c} car- 
bon dioxide? 

22. Describe the usual method of preparing carbon dioxide. Give 
the equation for the reaction. State its properties. 

23. Describe liquid and solid carbon dioxide. How are they pre- 
pared ? For what are they used ? 

24. What is the relation of carbon dioxide to animal and to plant 

25. State fully the relation of carbon dioxide to the unstable acid 
H.,CO 3 . Give the equations for the formation and decomposition of 
this acid. 

26. What are carbonates? Name three. How are they formed? 
What are their properties ? 

27. What is (a) " soda water," () carbonated water, (c) carbonic 
acid, (d) carbonic oxide, (e) carbonic anhydride, (/) limestone or 

28. What is the difference between (a} sodium carbonate and sodium 
bicarbonate, and (b) calcium carbonate and acid calcium carbonate ? 

29. Why is (a) CO 2 the formula of carbon dioxide, and (b} CO of 
carbon monoxide? 

30. State briefly the history of carbon dioxide. 

31. Give a brief account of (a} Black, (b) Van Helmont, and 
(c) Bergman. 

32. Illustrate the law of multiple proportions by the oxides of 

33. Give the equations for (a} the oxidation of carbon to carbon 
monoxide, (^) the reduction of carbon dioxide to carbon* rnonoA^-.. 

34. How is carbon monoxide (#) formed, and (<) usually prepared ? 

35. What is the relation of carbon monoxide to water gas? 

36. What are the properties of carbon monoxide? 

37. Illustrate Gay-Lussac's law by the combustion of carbon mo- 
noxide (2 CO + O., = 2 CO 2 ) . 

38. Illuminating gas, water gas, and the gas which escapes from a 
coal fire are poisonous. Why? 

39. What is cyanogen? Hydrocyanic acid? Describe potassium 
cyanide. For what is it used? Describe ootassium sulphocyanate. 
State its chief use. 

Carbon and its Oxides. 101 

40. The specific gravity of charcoal is about 1.5. Why does it float 
on water? 

41. How can carbon monoxide and carbon dioxide be changed into 
each other? 

42. Review (a) combustion, () solution of gases (especially carbon 
dioxide) in water, (c} respiration. 

43. State and explain the various chemical changes which occur from 
the entrance of oxygen (in the air) below the grate of a red-hot coal 
fire to the end of the burning of the carbon monoxide at the top of the 


1 . How many grams of calcium carbonate are needed to prepare 
132 gm. of carbon dioxide ? 

2. What weight of carbon burned in air will produce n gm. of 
carbon dioxide ? 

3. Calculate the percentage composition of (a} calcium carbonate, 
() carbon monoxide, (c) carbon dioxide, (d) magnesium carbonate. 

4. What per cent of carbon (by weight) is contained in carbon 
monoxide and in carbon dioxide ? 

5. If 20 gm. of carbon are heated in the presence of 44 gm. of 
carbon dioxide, (a) what weight of carbon monoxide is formed, and (<) 
what weight, if any, of carbon remains ? 

6. How many liters of carbon dioxide must be passed over red-hot 
charcoal to yield 84 gm. of carbon monoxide ? 

7. How much carbon dioxide () by weight and () by volume is 
in the air of a room 6 m. long, 4 m. wide, and 3 m. high, if there is 
i vol. of carbon dioxide in 1000 vol. of air ? 

8. What weight of water must be decomposed to furnish enough 
oxygen to form (with pure carbon) 44 gm. of carbon dioxide ? 

9. How many grams of calcium carbonate will produce 15 1. of 
carbon dioxide ? 

10. If a piece of pure graphite weighing 7 gm. is burned in oxygen, 
what volume of carbon dioxide is formed ? 



Hydrocarbons are compounds of carbon and hydrogen. 
They number about two hundred, and their properties 
vary between wide limits. They are found in petroleum 
and its products (kerosene, naphtha, lubricating oils, par- 
affin wax, etc.), in coal tar, in coal gas and natural gas, 
and in some essential oils, such as turpentine. On a large 
scale they are prepared by the destructive distillation of 
petroleum, wood, coal, and coal tar. Indirectly the hydro- 
carbons are the source of many other compounds of car- 
bon, which are extensively used in numerous industries. 

The existence of so many hydrocarbons is due to the fact that atoms 
of carbon have power to unite with themselves. This property gives 
rise to compounds which form natural groups or series. Simple rela- 
tions exist between many hydrocarbons, especially between members 
of the same series. The consecutive members of a series differ in com- 
position by CH 2 . Thus, in the methane series, methane is CH 4 and 
ethane is C 2 H 6 ; in the ethylene series, ethylene is C 2 H 4 and propylene 
is C 3 H 6 ; in the acetylene series, acetylene is C 2 H 2 and allylene is 
C 3 H 4 ; and in the benzene series, benzene is C 6 H G and toluene is C-H 8 . 
These series are called homologous series. 

Methane is found in coal mines, being a gaseous prod- 
uct of the processes which changed vegetable matter 
into coal. It is called fire damp by miners. It is also 
formed in marshy places by the decay of vegetable matter 
under water, and is therefore often called marsh gas. 

Methane. 203 

It is a constituent of natural gas and petroleum, and forms 
a large proportion of the illuminating gas obtained by 
heating coal. 

Methane is usually prepared in the laboratory by heating a mixture 
of sodium acetate, sodium hydroxide, and quicklime in a hard glass or 
metal vessel, and collecting the gaseous product over water. It may 
also be prepared by the interaction of aluminium carbide and water, 

A1 3 C 4 +i2H 2 0= 3CH 4 + 4A1(OH) 8 

Aluminium Carbide Water Methane Aluminium Hydroxide 

Methane has no color, taste, or odor. It burns with a 
pale, luminous flame. A mixture of methane with oxygen 
or air explodes violently when ignited by a spark or flame. 
Terrible disasters occur in coal mines from this cause. The 
products of the explosion are carbon dioxide and water, 
thus- CR4 + 2 Q 2 = co? + 2 H 2 
Methane Oxygen Carbon Dioxide Water 

The carbon dioxide, called choke damp or black damp by 
the miners, often suffocates those who escape from the 

Other members of the methane series are ethane (C 2 H 6 ), propane 
(C 3 H 8 ), butane (C 4 H 10 ). This series is also called the paraffin series, 
on account of the chemical indifference of its members. It has the 
general formula C n H 2n + 2- Butane and the succeeding fifteen or 
twenty members are liquids, and the highest members are solids. 

Chlorine and hydrocarbons interact, that is, chlorine replaces hydro- 
gen, atom for atom. Thus 

CH 4 + 2C1 = CH 3 Cr + HC1 
Methane Chlormethane 

This chemical change is called substitution, and illustrates one of the 
methods used in preparing derivatives of carbon known as substitution 
products. The paraffins are saturated hydrocarbons. This means 
that the carbon in them is saturated, so to speak, with hydrogen, and 
has no tendency to unite directly with more atoms of hydrogen or 
other elements. 

204 Descriptive Chemistry. 

Ethylene or olefiant gas is formed by the destructive 
distillation of wood and coal. It is usually prepared by 
heating a mixture of concentrated sulphuric acid and ethyl 
alcohol, and collecting the gas over water. The alcohol 
decomposes into ethylene and water, the latter being ab- 
sorbed by the sulphuric acid. The essential change is 
represented thus 

C^WgO = C^H^ -f- H^O 
Alcohol Ethylene 

Ethylene is a colorless gas, and has a pleasant odor. It 
can be condensed to a liquid, which by evaporation pro- 
duces a temperature as low as 140 C. It burns with a 
bright, yellow flame, and is one of the illuminating constit- 
uents of coal gas. When ethylene burns, the complete 
combustion is represented thus 

C 2 H 4 + 3 2 = 2C0 2 + 2H 2 
Ethylene Carbon Dioxide Water 

If mixed with oxygen in this proportion and ignited, the 
mixture explodes. 

Other numbers of this series are propylene (C 3 H 6 ) and butylene 
(C 4 H 8 ). These are unsaturated hydrocarbons. Unlike the paraffins, 
they form addition products by uniting directly with other substances, 
especially chlorine, thus 

C 2 H 4 + C1 2 = C 2 H 4 C1 2 

Ethylene Ethylene Chloride 

Ethylene chloride is one of the two dichlorethanes ; they have the 
same percentage composition, molecular weight, and formula (C 2 H 4 C1 2 ), 
but are very different compounds. They illustrate isomerism and are 
called isomers. This kind of isomerism is called metamerism. The 
difference in properties is believed to be due to a different arrangement 
of the atoms in the molecules. Isomerism occurs frequently among 
carbon compounds. 

Acetylene. 205 

Acetylene is formed by the direct union of hydrogen 
and carbon when an electric arc is produced between two 
carbon rods in hydrogen gas. This method of formation, 
though not convenient, is interesting, because no other hy- 
drocarbon has as yet been directly built up from its elements. 
A small quantity is present in coal gas. It is also formed 
by the incomplete combustioft of coal gas, e.g. when the 
flame of a Bunsen burner strikes back and burns at the 
base (see Bunsen Burner). Acetylene is now prepared 
cheaply on a large scale by treating calcium carbide with 
water, thus 

CaC 2 + 2H 2 O = C 2 H 2 + Ca(OH) 2 

Calcium Carbide Acetylene 

Acetylene is a colorless gas, and, if impure, has an offen- 
sive odor. It is poisonous if breathed in large quantities, 
but much less dangerous than gases containing carbon 
monoxide. It is lighter than air, its density being about 
0.92. Water at the ordinary temperature dissolves its own 
volume of the gas. Reliable tests show that acetylene 
does not act upon any common metal or alloy, though it 
forms explosive compounds with salts of metals, especially 
copper. As a precaution, copper and brass are seldom 
used in large vessels containing or generating acetylene, 
though they might be safely used on small vessels like 
bicycle lamps. 

Under a pressure of 40 atmospheres and a temperature of 20 C. it 
liquefies. Cylinders of liquid acetylene have exploded, causing loss of 
life and destruction of property, and its use in this form has been pro- 
hibited in some localities. Under ordinary atmospheric conditions acety- 
lene will not explode. If compressed, it will explode when a spark or 
flame is brought near it. A mixture of acetylene and air, if ignited, 
explodes. The mixture to be explosive, however, must contain from 
about 3 to 65 per cent of acetylene (a condition hardly possible 


Descriptive Chemistry. 

except from sheer carelessness), because the disagreeable odor reveals 
the presence of the gas. Acetylene must be used with the same precau- 
tion as any other illuminating gas. 

Acetylene is found by analysis to contain only carbon and hydrogen 
combined in the ratio of 12 to I by weight. Its vapor density is 13. 
Therefore its molecular weight must be 26 and its formula C 2 H.,. 

Acetylene is an unsaturated hydrocarbon, and like ethylene combines 
directly with bromine, hydrogen, and other elements. When passed 
into silver or copper solutions, it forms explosive compounds called 
acetylides (e.g. Ag 2 C 2 and Cu 2 C 2 ). Heated to a high temperature, it 
changes into other hydrocarbons, one being benzene, thus 

3 C 2 H 2 = C 6 H 6 
Acetylene Benzene 

At a very high temperature (about 800 C.) it decomposes into carbon 
and hydrogen. The change of acetylene into benzene illustrates po- 
lymerism. Polymers have the same percentage composition, but 
different molecular weights (see Isomerism). 

Acetylene as an Illuminant. Acetylene burns in the 
air with a luminous, smoky flame. But when air is mixed 

with the gas as the latter 
issues from a small opening, 
the mixture burns with a 
brilliant, white flame, which 
does not smoke. It is grad- 
ually coming into use as an 
illuminant. The flame is 
almost like sunlight, hence 
by the acetylene flame most 
colors appear the same as in 
daylight. It is also adapted 
for taking photographs, since 
its action closely resembles 
that of the sun. It is a diffusive light, and the flame is 
much smaller than an ordinary gas flame of the same 
lighting power (Fig. 32). 

FlG. 32. Relative size of acetylene 
and illuminating gas flames giving the 
same amount of light. The acetylene 
(smaller) flame consumes only one 
tenth as much gas an hour as the illu- 
minating gas flame. (One half actual 

Petroleum. 207 

With a proper burner the combustion of acetylene is 
complete, and may be represented thus 

2C 2 H 2 + 5O 2 = 4CO 2 + 2 H 2 O 
Acetylene Oxygen Carbon Dioxide Water 

In most acetylene burners the gas issues from two 
small holes drilled at an angle, so that the jets strike 
each other and produce a flat flame 
(Fig. 33). Other holes, properly 
located, permit air to be drawn in 
mechanically by the acetylene as it 
rushes through the burner. The open- 
ings for the mixture are so fine that 
FIG. 33. Acety- the flame cannot strike back and cause FlG 34 ._ Acety- 
lene flame. an explosion (Fig. 34). lene burner. 

Generation of Acetylene. The ease with which acetylene is gener- 
ated can be shown by putting a little water in a test tube and then drop- 
ping in small lumps of calcium carbide. The gas bubbles through the 
liquid ; after the action has proceeded long enough to expel the air, 
the acetylene may be lighted by holding a burning match at the mouth 
of the tube. On a larger scale, the gas can be generated by putting the 
calcium carbide into a flask provided with a dropping funnel and de- 
livery tube, and allowing water to drop slowly upon the carbide ; the 
gas thus generated can be collected in bottles over water. There are 
two classes of commercial generators. In one, water is added to the 
calcium carbide, but in the other the carbide drops into the water. The 
intense heat liberated when calcium carbide interacts with water de- 
composes acetylene ; hence, a generator to be effective and safe should 
be constructed so that this heat will be absorbed. The first class of 
generators is dangerous, except when a small quantity of gas is desired, 
as on the lecture table or in a bicycle lantern. -In the second class, a 
small amount of calcium carbide drops automatically into a large vol- 
ume of water as fast as the gas is needed, thus insuring a pure, cool 
gas, and eliminating the danger of an explosion. A pound of calcium 
carbide yields about five cubic feet of acetylene gas. 

Petroleum is the source of many useful hydrocarbons. 
It is an oily liquid obtained from the earth in many parts 
of the world. In the United States the chief localities are 

208 Descriptive Chemistry. 

Ohio, New York, Pennsylvania, West Virginia, Kentucky, 
Indiana, Colorado, Texas, and California. The immense 
deposits in Russia are in the Baku district on the Caspian 
Sea. Some is also found in Canada, India, Japan, and 

Crude petroleum is a thick liquid, with an unpleasant 
odor. Its color varies from straw to greenish black, and 
most kinds are greenish in reflected light. It usually floats 
upon water. Its composition is complex, but all varieties 
are essentially mixtures of many hydrocarbons. Ameri- 
can oils contain chiefly members of the paraffin series. 
Some varieties contain compounds of nitrogen and of 

In some localities the oil issues from the earth, but it is usually neces- 
sary to drill through rocks and insert a pipe into the porous rock 
containing oil. At first the oil often "shoots' 1 out of the well in 
tremendous volumes, owing to the pressure of the confined gas, but 
after a time a pump is needed to draw it to the surface. The oil is then 
forced by powerful pumps through large pipes to central points for 
storage or for delivery to refineries, which are often many miles from 
the oil well. This network of pipes in the eastern United States is over 
25,000 miles long. 

Some crude petroleum is used in making water gas (see 
below), and as fuel on locomotives and steamships, but 
most of it is separated into various commercial products. 
This process, which also involves purification, is called re- 
fining. The petroleum is distilled in huge iron vessels, 
and the vapors are condensed as they pass through coiled 
pipes immersed in cold water. Certain products are ob- 
tained from the residue left in the still. 

The different distillates, which are collected in separate tanks, are 
further separated and purified by redistillation. The commercial 
products obtained from the first distillation are cymogene, rhigolene, 
gasolene, naphtha, benzine, and kerosene. These liquids are mixtures 

Natural Gas. 209 

of several different hydrocarbons. They are widely used as solvents, 
fuels, and in making gas. 

Kerosene is the well-known illuminating oil. Being the most valu- 
able product from petroleum, it is very carefully freed from inflammable 
liquids and gases, which might cause an explosion, and from tarry 
matter and semi-solid hydrocarbons, which would clog the wicks of 
lamps. This is done by agitating it successively with sulphuric acid, 
sodium hydroxide, and water. Commercial kerosene must have a legal 
flashing point. This is "the temperature at which the oil gives off 
sufficient vapor to form a momentary flash when a small flame is 
brought near its surface." In most states the flashing point is 44 C. 
(or iiiF.). 

From the residuum left in the still after the first distillation many 
grades of lubricating oil, vaseline, and paraffin wax are obtained 
by further treatment. Mineral lubricating oils have largely replaced 
animal and vegetable oils. Vaseline finds extensive use as an ointment. 
Paraffin wax is used to make candles, to water-proof paper, to extract 
oils from plants and flowers, and as a coating for many substances, 
thereby producing a smooth surface or facilitating slow combustion (as 
in parlor matches). The final residue is coke. Hydrocarbons are 
often extracted from it, some is made into electric light carbons, and 
some is used as a fuel. 

This vast industry yields over two hundred different commercial 
products, many of them being indispensable to the comfort and conven- 
ience of mankind. In 1901 the United States produced over 69,000,000 
barrels of crude petroleum. 

The Origin of Petroleum is doubtful. Some think it was produced 
by the decomposition or slow distillation of plants and animals. 
Recently it has been suggested that it resulted from the interaction of 
water and metallic carbides, especially iron carbide, at great depths. 

Natural Gas is a combustible gas, which issues from the 
earth in many places. Methane is the principal constituent 
of the mixture. It is used as a fuel for heating houses, 
generating steam, and manufacturing iron, steel, glass, 
brick, and pottery. 

In Ohio, Indiana, and other gas-producing regions of the United 
States, wells, like petroleum wells, are drilled for the escape of natural 

2io Descriptive Chemistry. 

gas, which is distributed to consumers through pipes similar to those 
used for illuminating gas. Enormous quantities are consumed in the 
United States, the annual product being valued at over $20,000,000. 

Illuminating Gas. Besides acetylene there are other 
kinds of illuminating gas. Coal gas and water gas are the 
most common. 

Coal Gas is made by distilling bituminous coal and puri- 
fying the volatile product. The hydrogen in the coal 
passes off partly as free hydrogen, and partly in combina- 
tion with carbon as hydrocarbons, and with nitrogen as 
ammonia. The ammonia, carbon dioxide, and sulphur 
compounds are regarded as impurities, and are removed 
before the gas is sent to the consumer. The essential 
parts of a coal-gas plant are shown in Figure 35. 

The coal is distilled in a -shaped retorts, made of fire clay and 
about eight feet long. Six or more retorts are arranged in tiers form- 
ing a group or bench, so that all the retorts of a bench can be heated 
by a single fire usually of coke. Several benches placed end to end 
constitute a stack. The retorts are heated red hot, and about two hun- 
dred pounds of coal are evenly distributed on the bottom of each retort 
with a long iron scoop, and the mouth is quickly and tightly closed by 
an iron lid. The distillation continues from four to six hours, during 
which the temperature often reaches 1200 C. The lid is then removed, 
the red-hot coke is pushed or raked out, and another charge of coal is 
quickly introduced. The coke is quenched with water to prevent fur- 
ther combustion. Some of it is used for heating the retorts, but a part 
is sold. 

The volatile products pass from each retort up through a standpipe, 
down the dip pipe, and bubble through water into the hydraulic main. 
This is a horizontal, half-round pipe extending the whole length of the 
stack. Here some of the tar is deposited and ammonium compounds 
are dissolved by the water which flows constantly through the main. 
This water is kept at the same level and acts as a " seal " to prevent the 
gas from passing back into the retorts. The ammoniacal liquor and 
tar flow into a tar well. 

From the hydraulic main the gas which is hot and impure passes 

Illuminating Gas. 



212 Descriptive Chemistry. 

into the condenser. This is a series of vertical iron pipes, several 
hundred feet long. They are connected at the top, but they open at 
the bottom into a series of boxes so constructed that the gas must pass 
through the entire length of the pipes, while the tar and ammoniacal 
liquor flow into 'the tar well. The main object of the condenser is to 
cool the gas slowly and condense and remove the tar. 

An exhauster, in most plants, draws or forces the gas from the 
hydraulic main through the condenser into the scrubber and onward 
through the purifiers into the gas holder. The exhauster also reduces 
the pressure in the retorts and regulates the pressure in the holder (see 
below) . 

The scrubber is a washing machine. Its purpose is to remove the 
remaining ammonia, part of the carbon dioxide, and hydrogen sulphide 
gas, and the last traces of tar. Scrubbers vary in construction. One 
form is a double tower filled with wooden slats or with trays covered 
with coke or pebbles over which ammoniacal liquor slowly trickles in 
the first part and pure water in the second. The gas enters at the 
bottom, meets the descending liquid, and is thoroughly washed. 
Another form widely used consists of a cylindrical vessel in which 
numerous wooden slats revolve in compartments and dip into am- 
moniacal liquor or water at the bottom. The liquid forms a film on 
the slats and absorbs the ammonia and other gases, while the resulting 
solution mixes with liquor at the bottom and flows into the proper well. 
Sometimes a separate tar extractor is connected with the scrubber. 
This is a tower filled with perforated plates, which catch and remove 

the tar mechanically as the gas 
passes through into the scrubber. 

From the scrubber the gas 
passes into the purifiers. Their 
FIG. 36. -Slat frame (or grid) used in chief purpose is to remove the 
the lime purifier. remaining carbon dioxide and sul- 

phur compounds. They are shal- 
low, rectangular iron boxes provided with slat frames loosely covered 
with lime (Fig. 36). In some plants iron oxide is used as the purifying 

The purified gas next passes through a large meter, which records 
its volume, into a gas holder. The holder is an enormous, cylindrical, 
iron tank in which the gas is stored. It floats in a cistern of water, and 
rises or falls as the gas enters or leaves. Weights and the pressure 

Water Gas. 213 

from the exhauster so balance it that it exerts just enough pressure to 
force the gas through the pipes to the consumer. 

A ton of good coal yields about 10,000 cubic feet of gas, 1400 pounds 
of coke, 120 pounds of tar, 20 gallons of ammoniacal liquor, and a vary- 
ing amount of gas carbon. The coke is a valuable fuel and finds a 
ready sale. The tar, or coal tar as it is often called, collected from the 
hydraulic main and condenser, is a thick, black, foul-smelling liquid. 
It was formerly thrown away. Some is used for preserving timber, 
making tarred paper and concrete, and as a protective paint. Most of 
it is now separated by distillation into its more important constituents, 
especially benzene (C 6 H C ) . These carbon compounds and their numer- 
ous derivatives appear in commerce as oils, medicines, dyestufFs, flavors, 
perfumes, and other useful products. The ammoniacal liquor from the 
hydraulic main, condenser, and scrubber is the source of ammonia and 
its compounds. Gas carbon is the hard deposit which collects on the 
inside of the retort, and is used in the electrical industries (see Gas 
Carbon). The sale of these by-products reduces the cost of making 
the coal gas. 

Water Gas is made by forcing steam through a mass of 
red-hot coal and mixing the gaseous product with hot gases 
obtained from oil. The essential parts of the apparatus 
are shown in Figure 37. 

Air is forced through the coal fire in the generator, and the hot 
gases which are produced pass down the carburetor, up into the super- 
heater, and escape through its top into the open air. This operation 
lasts about four minutes, and is called the " blow." It heats the fire 
brick inside the carburetor and superheater intensely hot, air often being 
forced in to raise the temperature. The air valves and the top of the 
superheater are now closed, and the " run " begins, which lasts about 
six minutes. Steam is forced into the generator at the bottom. In 
passing through the mass of incandescent carbon the steam and carbon 
interact thus 

C + H 2 O CO + H 2 

Carbon Steam Carbon Monoxide Hydrogen 

This mixture of hydrogen and carbon monoxide burns with a feeble 
flame, and before it can be used as an illuminating gas it must be 


Descriptive Chemistry. 

Characteristics of Illuminating Gases. 215 

enriched with gases which are illuminants. Therefore, the mixed gases 
pass to the top of the carburetor, where they meet a spray of oil. And 
as the gaseous mixture passes down the carburetor and up the super- 
heater, the hydrocarbons of the oil are transformed by the intense heat 
into hydrocarbons that do not liquefy when the gas is cooled. The ad- 
dition of hydrocarbons is called carbureting. From the superheater 
the water gas passes through the purifying apparatus into a holder. 

Water gas is seldom burned alone, but is usually mixed 
with 60 or 70 per cent of coal gas. This mixture is popu- 
larly called " illuminating gas." Owing to the high percen- 
tage of carbon monoxide, water gas and gases containing 
it are poisonous. 

Characteristics of Illuminating Gases. Both coal gas 
and water gas have a disagreeable odor. They are mix- 
tures having a composition which varies with the coal 
used, the temperature reached, and the degree of purifica- 
tion attained. The following table shows the average 





Marsh 2fas . . 


19 8 

Ethylene (and other illuminants) 



52 I 

Carbon monoxide ... 

7 2 


26 I 

Carbon dioxide 

I i 





Both kinds of illuminating gas may contain a little oxygen, and 
traces of ammonia and hydrogen sulphide gases. Nitrogen and the 
last portions of carbon dioxide are impurities not easily removed. 
Marsh gas, hydrogen, and carbon monoxide burn with a feeble (non- 
yellow) flame, and are often called diluents ; they furnish heat, but no 

216 Descriptive Chemistry. 

The luminosity of illuminating gas depends mainly 
upon the presence of hydrocarbons containing a relatively 
large proportion of carbon. Acetylene gas, which gives 
such a brilliant light, consists almost wholly of this hydro- 
carbon containing 90 per cent of carbon. The most im- 
portant illuminants in coal gas and water gas are ethylene 
and similar hydrocarbons, acetylene, and benzene (C 6 H 6 ). 

The commercial value of an illuminating gas depends upon its illu- 
minating power. This property is measured by a photometer and is 
expressed in * candles." The determination is made by comparing the 
light produced by burning the gas in a standard burner at the rate of 
five cubic feet an hour with the light produced by a standard wax candle 
burning at the rate of 120 grains (7.77 gm.) an hour. If the gas flame 
is 20 times brighter than the candle flame, then the candle power of the 
gas is 20. The candle power of ordinary coal gas is about 17, and 
that of water gas is about 25. Ordinary illuminating gas has a candle 
power of about 20, since it is usually a mixture of coal gas and water 

Flame. A flame is a mass of burning gas. Ordinarily 
it is gas combining chemically with the oxygen of the air. 
In the illuminating gas flame the gas itself is burning in 
the air. In a lamp flame the gas which burns comes from 
the oil which is drawn up the wick by capillary attraction, 
and then volatilized by the heat. Similarly, in a candle 
flame the burning gas comes from the melted wax. The 
flame produced by most burning hydrocarbons is yellowish 

The hydrocarbon flame has several distinct parts, though 
the structure of the flame is essentially the same, whether 
produced by burning illuminating gas, kerosene oil, or can- 
dle wax. The candle flame may be taken as the type. An 
examination of the enlarged vertical section shown in Fig- 
ure 38 reveals four somewhat conical portions, (i) Around 
the wick there is a black cone (A), filled with combustible 



FIG. 38. Candle 

gases formed from the melted wax. They do not burn be- 
cause no oxygen is present. With a glass tube of fine 
bore it is possible to draw off these gases 
from a large flame and light them at the 
upper end of the tube. (2) Around the 
lower part of the dark cone is a faint bluish 
cup-shaped part (>, B). It is the lower por- 
tion of the exterior cone where complete 
combustion of the gases occurs, since plenty 
of oxygen from the air reaches this portion. 
(3) Above the dark cone is the luminous 
portion (C). It is the largest and most 
important part of the flame. It is popu- 
larly spoken of as " the flame." Combus- 
tion is incomplete here, because little or no 
oxygen can pass through the exterior cone. The tempera- 
ture is high, however, and the hydrocarbons undergo 
complex changes. Acetylene is probably formed. The 
most characteristic change is the liberation of small par- 
ticles of carbon. This liberated carbon heated to incan- 
descence by the burning gases makes the flame luminous. 
The carbon glows but does not burn up, because little or 
no oxygen is present. A crayon or glass rod held in this 
part of the flame is at once coated with soot, which consists 
of fine particles of carbon. The exterior cone (D, D) is 
almost invisible. Here combustion is complete, because 
the oxygen of the air changes all the carbon into carbon 
dioxide. That this is the hottest region 
may be easily shown by pressing a piece of 
stiff white paper for an instant down upon 
the flame almost to the wick. The paper 
FIG. 39 .- Paper wi ^ fr Q charred by the outer part of the 

charred by a can- 
dle flame. flame, as shown in Figure 39. 

2i 8 Descriptive Chemistry. 

These four portions may be found in all luminous hydrocarbon 
flames, whatever the shape. An ordinary gas flame is flattened by forc- 
ing the gas flame through a narrow slit in the burner, so that the flame 
will give more light. The blue part is easily seen, however, when the 
gas flame is turned low or looked at through a small opening ; the dark 
and yellow parts are always visible the latter being intentionally en- 
larged. The flat or circular flame of an oil lamp likewise presents the 
same characteristics. 

The gaseous products of the combustion of hydrocarbons 
are water vapor and carbon dioxide. A bottle in which a 
candle is burning has, at first, a deposit of moisture on the 
inside ; and if the candle is removed and limewater added, 
the presence of carbon dioxide is shown by the milkiness of 
the limewater. The oxygen needed by the burning hydro- 
carbons is obtained from the air. If not enough oxygen is 
present, the flame smokes, i.e. the carbon is thrown off into 
the air before the particles are heated hot enough to glow. 
All oil lamps are so constructed that air enters the burner 
below the flame. Large oil lamps have a central opening 
through which a large volume of air passes up inside the 
circular flame. Otherwise the lamp would burn with a 
very smoky flame. 

The luminosity of hydrocarbon flames is affected by other things 
besides the presence of glowing carbon. One of these is temperature. 
Gases cooled before being burned give poor light. A candle flame may 
be cooled enough to extinguish it. Thus, if a coil of copper wire is 
lowered upon a candle flame, the flame smokes, loses its yellow color, 
and finally goes out ; but if a coil of hot wire is used, the flame burns 
unchanged. Gases, as well as solids and liquids, have a kindling tem- 
perature, i.e. a temperature to which they must be heated before they 
" catch fire." This temperature differs with different substances. As 
we lower the temperature of gases burning with a luminous flame, their 
luminosity decreases, and below their kindling point they will not burn. 
The density of the gases in the flame and of the atmosphere itself like- 
wise modifies luminosity. A candle flame was found by experiment to 
be smaller on the top of Mont Blanc than at the base. 

The Bunsen Burner and its Flame. 219 


Not all flames are luminous. The hydrogen flame is almost invisible, 
and the flames of carbon monoxide and methane are a faint blue. These 
flames yield no solid particles of carbon, but only gaseous products. The 
most common non-luminous flame is the Bunsen flame. 

The Bunsen Burner and its Flame. When illuminat- 
ing gas is mixed with air before burning, and the mixture 
burned in a suitable burner, a flame is produced which is 
non-luminous -and very hot. The 
temperature of the hottest part is 
about 1 500 C. This flame deposits 
no carbon, since its products are 
entirely gaseous. Such a flame is 
called the Bunsen flame, because 
it is produced in a burner devised 
by the German chemist Bunsen. 
This burner is constantly used in 
chemical laboratories as a source 
of heat, and modified forms have 
numerous uses. One form, for 
example, furnishes the heat in the 
gas range used for cooking. The 
parts of an ordinary Bunsen burner 
are shown in Figure 40. The gas 
enters the base and escapes through 
a very small opening into the long 
tube, which screws down upon this 
opening. At the lower end of the 
long tube there are two holes, 
through which air is drawn by the gas as it rushes out of 
the small opening. The gas and air mix as they rise in the 
tube, and this mixture of air and gas burns at the top of 
the long tube. The size of the air holes at the bottom of 
the long tube may be changed by a movable ring, thus 

FIG. 40. Parts of a Bunsen 

220 Descriptive Chemistry. 

varying the volume of the entering air. When the holes 
are open, the typical colorless, hot Bunsen flame is formed. 
The combustion of the hydrocarbons is practically com- 
plete. They burn up before particles of carbon are 
liberated, thus making the flame non-luminous and free 
from soot. Apparatus heated by this flame is not black- 
ened. The Bunsen flame may be made momentarily 
luminous by shaking or blowing fine particles into the 
flame, such as powdered charcoal dust, finely divided 
metals, and sodium compounds. 

It was formerly believed that the non-luminous character of the 
Bunsen flame is solely due to the complete combustion of the carbon by 
the oxygen of the entering air. Recent experiments have shown, how- 
ever, that the result is partly due to the diluting action of the nitrogen* 
The gas burns at top of the tube and not inside, because the proper 
mixture of gas and air flows out more quickly than the flame can travel 
back. If the gas supply is slowly decreased, the flame becomes smaller 
and finally disappears with a slight explosion. This change is called 
"striking back." It is due to the fact that the tube contains an explo- 
sive mixture of air and illuminating gas, through which the flame travels 
faster than the mixture escapes from the tube. This explosion illus- 
trates in a small way what often happens when a mixture of air and 
illuminating gas is ignited. Sometimes the flame is not extinguished, 
but burns within (and sometimes without) the tube. This flame has a 
pale color, a disagreeable odor, and deposits soot. 

The Bunsen flame has many characteristic properties. 
Its color is bluish, and the different corres have different 
colors. There are really three cones: (i) the blue or 
greenish inner one of unburned gases ; (2) the very faint 
blue middle one ; (3) and the outer one, which is pale blue, 
and represents the blue cone in the candle flame. The 
middle and outer cones are not always easily distinguished ; 
and for all practical purposes it is convenient to divide the 
flame into two parts, an inner cone of unburned gases 

Oxidizing and Reducing Flames. 221 

FIG. 41. The effects of wire gauze on a 
Bunsen flame. 

and an outer cone in which all the carbon is consumed. 
Combustible gases may be drawn off by a tube from the 
inner cone and ignited. A match laid for an instant across 
the top of the tube is 
charred only at the two 
points where it touches 
the outer cone ; and a 
sulphur match'- suspended 
by a pin across the top 
of an unlighted burner 
is not kindled when the 
gas is first lighted. A 
piece of wire gauze pressed down upon the flame shows 
a dark central portion surrounded by a luminous ring. 
The flame is beneath the gauze, although the gas passes 
freely through it and escapes. If the gas is extinguished 
and then relighted above the gauze, it will burn above 
but not beneath (Fig. 41). The gauze cools the gas below 
its kindling temperature. 

The miner f s safety lamp invented by Davy depends 
upon this last principle. It is an oil lamp surrounded 
by a cylinder of fine wire gauze (Fig. 42). When 
taken into a mine where there are explosive gases (fire 
damp), the flame continues to burn inside, though its 
size and color change. The gas often enters the lamp 
and burns inside, but the flame within does not ignite 
the gases without because the wire gauze keeps them 
cooled below their kindling temperature. Hence an 
explosion is often prevented. When miners notice 
changes in the lamp flame, they usually seek a safe 
FIG. 42. One place, 
form of Davy's * 

safety lamp. Oxidizing and Reducing Flames. The 

outer portiqn of the Bunsen flame is called the oxidizing 
flame, because here the oxygen is freely given to sub- 


Descriptive Chemistry. 

-~Y- A 

stances. The inner portion is called the reducing flame, 
because here the hydrocarbons withdraw 
oxygen. A sketch of the general relation 
of these flames is shown in Figure 43. A 
is the most effective part of the oxidizing 
flame, and B of the reducing flame. At 
A metals are oxidized, and at B oxygen 
compounds are reduced. 

Sometimes a long tube with a small opening 
at one end, called a blowpipe, is used to produce 
these flames. A tube with a flattened top is put 
inside the burner tube to produce a luminous flame. 
The tip of the blowpipe rests 
in or near this flame, and if 
air is gently and continuously 
blown through the blowpipe, 
a long, slender flame is pro- 
duced, called a blowpipe 
flame (Fig. 44). It is like 
the Bunsen flame as far as 
its oxidizing and reducing 

properties are concerned. The blowpipe is used 
in the laboratory and by jewelers and mineral- 
ogists. On a large scale the blowpipe flame is used to reduce or oxidize 
ores and to melt refractory substances (see Compound Blowpipe) . 

The Bunsen flame has recently been utilized in producing the Wels- 
bach light. The non-luminous flame heats an inverted bag or " man- 
tle " of oxides of rare metals, and the mantle glows with an intense 
light. The candle power varies from 40 to 100. This form of burner 
is widely used because it produces a brilliant light. 


1. What are hydrocarbons ? ^ Where are they found ? Name sev- 
eral familiar substances containing hydrocarbons. 

2. Are there many hydrocarbons ? Why ? 

3. What is an homologous series of hydrocarbons ? Name four 
such series. 

FIG. 43. The oxi- 
dizing (A) and reduc- 
ing (Z?) flames. 

FIG. 44. Blowpipe 
flame, showing oxidiz- 
ing (A) and reducing 
(B) parts. 

Exercises. 223 

4. What is methane ? What other names has it ? Where is it 
found ? How is it usually prepared ? State its essential properties. 
Why is it a dangerous gas ? Illustrate your answer by an equation. 

5. What other name has the methane series ? Why ? Illustrate 
the following terms by the paraffin series : (a) substitution, () substi- 
tution product, (6-) saturated hydrocarbon. 

6. What is ethylene ? How is it prepared ? Where is it found ? 
State its properties. Give the equation expressing the combustion of 

7. Illustrate the following terms by the ethylene series : (a) unsatu- 
rated hydrocarbon, (b) addition product, (c) isomerism, (d} metamer- 
ism, (e) isomer. 

8. Review the subject of calcium carbide (see Chapter X). 

9. What is acetylene ? How is it formed ? How is it prepared ? 
Give the equation for the reaction. Summarize the properties of acety- 

10. Illustrate the following terms by acetylene : (a) polymerism, 
(^) polymer, (V) unsaturated hydrocarbon. 

1 1 . Describe the acetylene (a) flame, () burner, and (c) generator. 
What precautions must be observed in using acetylene as an illuminant ? 

12. What is (a) choke damp, () black damp, (V) marsh gas, 
(//) olefiant gas ? 

13. What is the formula of (a) methane, () ethylene, (c) benzene ? 
Why is C 2 H 2 the formula of acetylene 4 ? 

14. How many volumes of oxygen are needed for the combustion of 
one volume of (a) methane, () ethylene, and (c) of two volumes of 
acetylene ? What volumes of what products are formed in each case ? 
What law do these relations illustrate ? 

15. What is petroleum ? Where is it found ? Of what is petroleum 
composed ? How is it obtained from the earth ? Describe briefly the 
refining of petroleum. 

1 6. What is kerosene ? Describe its method of preparation. Define 
and illustrate the \ES\bfldskingpoint. 

17. State the uses of (a) gasoline, (<) lubricating oils, (c) vaseline, 
(tf) paraffin wax. 

1 8. What is natural gas ? Where is it found ? Of what is it com- 
posed ? For what is it used ? 

19. What is coal gas ? Describe briefly its manufacture. 

20. What is coal tar ? What are its uses ? 

224 Descriptive Chemistry. 

21. What is ammoniacal liquor ? What is its source ? How is it 
obtained ? For what is it used ? 

22. Review (a) coke, and (b} gas carbon (see Chapter XIV). 

23. What is water gas ? Describe briefly its manufacture. What is 
meant by " enriching " water gas ? What is producer gas ? 

24. Give the equation for the interaction of carbon and steam. How 
many volumes of steam are needed to produce one volume of each of the 
products ? 

25. What is illuminating gas ? State its chief properties. What 
are its (a} light-giving constituents, (b} diluents, (c) impurities ? Upon 
what does its luminosity depend ? How is this property measured and 
expressed ? Give two reasons why illuminating gas is dangerous. 

26. What is a flame ? Illustrate your answer. Describe the struc- 
ture of a candle flame. What are the chief gaseous products of combus- 
tion ? Why do lamps sometimes smoke ? What affects the luminosity 
of many flames ? 

27. Describe (a) the Bunsen flame, (fr) the Bunsen burner. Why 
is the Bunsen flame non-luminous ? Describe and explain the " strik- 
ing back" of the Bunsen flame. Describe the structure of the Bunsen 
flame. What is the miner's safety lamp, and upon what principle is it 
constructed ? 

28. Review oxidation and reduction. 

29. What is (#) an oxidizing flame ? Describe a blowpipe and its 
flame. For what is it used ? 

30. Describe the Welsbach light. 


1. Calculate the percentage composition of (a) methane (CH 4 ), 
() ethylene (C 2 H 4 ), and (c} acetylene (C 2 H 2 ). 

2. What weight of oxygen is needed for the complete combustion 
of 4 gm. of ethylene ? (Equation is C 2 H 4 + 3<3 2 = 2 CO, + 2 H 2 O.) 

3. What is the simplest formula of a compound having the compo- 
sition H = 7.69 and C = 92.3 ? 

4. Calculate the molecular formula of a compound having the vapor 
density 38.8 and the composition C = 92.3 and H = 7.69. 


FLUORINE, bromine, and iodine, together with chlorine, 
are often grouped, and called the fialogenj. They resem- 
ble each other in a general way, aiuT forni analogous com- 
pounds which have similar properties, differing mainly in 

Halogen means " a sea-salt producer." It is applied to this group 
of elements because they form salts which resemble sodium chloride 
(common sslt or sea salt). Chlorides, bromides, and iodides are some- 
times called haloid salts or halides. The Greek word for salt, hals, 
suggested these terms. 


Occurrence. Fluorine is the most active of all the ele- 
ments, and is therefore never found free in nature. It 
occurs abundantly in combination with calcium as fluor 
spar or calcium fluoride (CaF 2 ). Other native compounds 
are cryolite (Na 3 AlF 6 ) and apatite (CaF 2 . 3 Ca 3 (PO 4 ) 2 ). 
Minute quantities of combined fluorine are found in bones 
and blood, in the enamel of the teeth, and in sea and some 
mineral waters. 

Fluorine is named from fluor spar, which melts easily and is used as 
a flux to make substances flow together (hence the derivation from the 
Latin fluo, I flow). 

The Isolation of Fluorine was accomplished in 1886 by 
Moissan, though many unsuccessful attempts had been 
previously made. He decomposed hydrofluoric acid by 



Descriptive Chemistry. 

electricity and collected the liberated fluorine. The 
achievement was attended with tremendous difficulties, 
owing to the intense activity of fluorine and its corrosive 

The essential parts of the apparatus used by Moissan are shown in 
Figure 45. The U-tube, made of an alloy of platinum and iridium, is 
provided with tightly fitting stoppers of fluor 
sr (S. S) . Through the stoppers pass the elec- 
trodes (E, E) of platinum iridium, held in place 
by screw caps (C,C). Side tubes ( T, T) allow 
the liberated gases (fluorine and hydrogen) 
to be drawn off separately through platinum 
delivery tubes. Perfectly dry hydrofluoric 
acid is put into the U-tube and dry acid 
potassium fluoride (HKF 2 ) is added to enable 
the solution to conduct the current liquid 
hydrofluoric acid itself being a non-conductor. 
The U-tube is cooled to a very low tempera- 
ture (23 to 50 C.), and on passing a 
current through the apparatus fluorine is 
evolved at the positive electrode and hydrogen 
at the other. The fluorine, freed from hydro- 
fluoric acid vapor, was collected by Moissan 
at first in a platinum tube with thin fluor spar plates closing each end, 
so that he could look inside and examine the gas. Later he found that 
pure fluorine can be collected in glass tubes, since it attacks glass only 
very slowly. 

Properties. Fluorine has a sharp odor and a greenish 
yellow color, but lighter and more yellowish than chlorine. 
Its density is 1.265 ( an " = 0- Subjected to pressure and 
a very low temperature, it condenses to a pale yellow liquid, 
which boils at 187 C. The pure gas can be liquefied 
in a glass vessel. Chemically, fluorine is intensely active. 
Hydrogen, bromine, iodine, sulphur, phosphorus, carbon, 
silicon, and boron take fire in it. Oxygen, nitrogen, and 
argon do not unite with it. Most metals burn in it, form- 

FlG. 45. Moissan's ap- 
paratus for preparing flu- 

Fluorine Bromine Iodine. 227 

ing fluorides. Gold and platinum are not attacked by it 
below red heat. Copper becomes coated with copper fluor- 
ide, which protects the metal, so that copper vessels may 
be used as fluorine generators. Moissan used a copper 
U-tube to prepare large volumes. Water is decomposed 
by it at ordinary temperatures, owing to the intense attrac- 
tion between hydrogen and fluorine ; hydrocarbons, for a 
similar reason, are instantly decomposed, hydrofluoric acid 
and carbon fluorides being the products. 

The exhaustive work of Moissan shows that fluorine, though more 
active than the other halogens, is similar to them, and should be regarded 
as the first member of that group. 

Hydrofluoric Acid, HF, is the compound of fluorine 
corresponding to hydrochloric acid. It is prepared by 
the interaction of a fluoride and concentrated sulphuric 
acid. Calcium fluoride is usually used, and the experi- 
ment is performed in a lead dish. The chemical change 
is represented thus 

CaF 2 + H 2 SO 4 = 2HF + CaSO 4 

Calcium Fluoride Sulphuric Acid Hydrofluoric Acid Calcium Sulphate 

Hydrofluoric acid, like hydrochloric acid, is a colorless 
gas, which fumes in the air and dissolves in water, the 
solution being the commercial hydrofluoric acid. Both 
gas and liquid are dangerous substances. The gas 
is extremely poisonous, and the liquid, if dropped on 
the skin, produces terrible sores. Owing to its corro- 
sive action the acid is preserved and sold in platinum, 
rubber, or wax bottles. The acid and the moist gas attack 
glass, and are used extensively in etching. The glass is 
coated with wax, and the design to be etched is scratched 
through the wax. The glass is the*n exposed to the gas or 
the liquid, which attacks the exposed places. When the 

228 Descriptive Chemistry. 

wax is removed, a permanent etching like the design is 
visible. Glass is an artificial compound of silicon a 
silicate. The corrosive action of hydrofluoric acid upon 
glass is due to the ease with which the acid decomposes 
glass and forms with the silicon a volatile compound, 
called silicon tetrafluoride (SiF 4 ). Since silicon dioxide 
(or sand) is the essential constituent of the mixture from 
which glass is made, the equation for etching glass may 
be written thus 

SiO 2 + 4HF = SiF 4 + 2 H 2 O 
Silicon Hydrofluoric Silicon 

Dioxide Acid Tetrafluoride 

Scales on thermometers and on other graduated glass 
instruments are etched with hydrofluoric acid. 

The vapor density of hydrofluoric acid gas indicates that its formula 
is HF at high temperature, but H 2 F 2 at lower temperatures (30 C.). 


Occurrence. Bromine is never found free in nature on 
account of its chemical "activity. Bromides are widely 
distributed, especially magnesium bromide. The salt 
springs of Ohio, West Virginia, Pennsylvania, and Michi- 
gan, and the salt deposits at Stassfurt in Germany furnish 
the main supply of the element. Sea water, Chili salt- 
peter (NaNO 3 ), and certain seaweeds contain a small 
quantity of combined bromine. 

Preparation. Bromine is obtained from its compounds 
by treatment with chlorine, or with sulphuric acid and 
manganese dioxide. In the laboratory, bromine is pre- 
pared by heating potassium bromide with manganese 
dioxide and sulphuric a'cid in a glass vessel. The bromine 
is easily liberated as a dense, brown vapor, which often 

Fluorine Bromine Iodine. 229 

condenses to a liquid and runs down the walls of the 
vessel. The chemical change is represented thus - 

2 KBr + 2 H 2 SO 4 + MnO 2 = Br 2 + MnSO 4 + K 2 SO 4 + 2 H 2 O 

Potassium Sulphuric Manganese Bro- Manganese Potassium Water 
Bromide Acid Dioxide mine Sulphate Sulphate 

Bromine is sometimes prepared by treating a bromide with 
manganese dioxide and hydrochloric acid. 

The source of commercial bromine in the United States is " bittern " 
a concentrated liquid left after salt is crystallized from brine. In the 
continuous process the hot bittern flows down a large tower filled with 
broken brick or burned clay balls ; chlorine gas and steam forced in at 
the bottom meet the bittern and liberate the bromine, which passes as a 
vapor out of the top into a condenser. The main chemical change is 
represented thus 

MgBr 2 + C1 2 - Br 2 -f MgCl 2 

Magnesium Bromide Chlorine Bromine Magnesium Chloride 

In the periodic process, used chiefly in the United States, a huge stone 
still is charged with manganese dioxide, hot bittern, and sulphuric acid, 
and heated by steam. The bromine distills into a condenser, as in the 
other process. Sometimes potassium chlorate is used as the oxidizing 

Properties.VtJ^romineis a heavy, reddish brown liquid 
at the ordinary T^njgCIamel Its specific gravity is about 
three. It is a volatile liqu!?f, boiling at about 59 C. The 
vapor, which is given off freely, has a disagreeable, suffo- 
cating odor. This property suggested the name bromine 
(from the Greek word bromos, a stench). It is poisonous, 
and burns the flesh frightfully. Bromine is somewhat 
soluble in water. The solution, called bromine water, has 
a brown color, and when cooled deposits a crystalline 
hydrate (Br 2 . 10 H 2 O). Many other properties of bromine 
are similar to those of chlorine. Thus, it combines with 
metals and other elements ; it also bleaches. 

230 Descriptive Chemistry. 

Compounds of Bromine are similar to those of chlorine. Hydrobro- 
mic acid (HBr) is a colorless, pungent gas, which fumes in the air and 
dissolves freely in water, forming the solution usually called hydrobromic ' 
acid. Its other properties closely resemble those of hydrochloric acid. 
Bromides are salts of hydrobromic acid, though many are formed by 
direct combination with bromine. Like the chlorides, most bromides 
dissolve in water. Potassium bromide (KBr) is a white solid, made by 
decomposing iron bromide with potassium carbonate. It is used exten- 
sively as a medicine and in photography (in preparing silver bromide 
plates and films). Bromides of sodium, ammonium, and cadmium have 
a limited use. 

Miscellaneous. Bromine itself is used to make potassium bromide 
and other compounds, especially a class of coal tar dyes used to color pink 
string and to make red ink. Annually over 500,000 pounds of bromine 
are prepared in the United States, while Germany exports about 400,000 
pounds of bromine, and 500,000 pounds of bromine compounds. 

Balard discovered bromine in 1826 in the mother liquor (or bittern) 
from brine. Liebig supposed it was chloride of iodine, and thus failed 
to discover it, because, as he said, he yielded to " explanations not 
founded on experiment." 


Occurrence. Free iodine is never found in nature, but 
like chlorine and bromine it is combined with metals, 
especially sodium, potassium, or magnesium. It is widely 
distributed, though the quantity in any one place is small. 
Tobacco, water cress, cod-liver oil, oysters, and sponges con- 
tain minute quantities. Native iodides of silver and of mer- 
cury are found. The ash of some seaweeds contains from 
0.5 to 1.5 per cent of its weight of iodides of sodium and 
potassium. Sodium iodate occurs in the deposits of salt- 
peter in Chili, and is now the main source of the element. 

Preparation. Iodine is prepared in the laboratory by a 
method similar to that used for bromine. Potassium iodide, 
manganese dioxide, and sulphuric acid are heated in a glass 
vessel, and the iodine appears as a violet vapor, which con- 

A . 

Fluorine Bromine Iodine. 231 

denses on the upper part of the vessel into dark grayish 

On a commercial scale iodine is prepared from the ash of seaweeds 
and from the mother liquors of Chili saltpeter, (i) Along the coasts 
of France, Scotland, and Norway seaweed is collected and burned, 
usually in closed vessels. The ash is called kelp or varec. The solu- 
ble portions are removed by agitation with water. The 'filtered liquid 
is further purified, and from the final mother liquor in which the iodides 
are dissolved, the iodine is extracted by heating with sulphuric acid and 
manganese dioxide. Sometimes chlorine is used to extract the iodine. 
In either case the mother liquor and its added ingredients are distilled 

FiG. 46. Apparatus for purifying iodine. 

gently in an iron pot with a lead cover, which is connected with two 
rows of bottle-shaped condensers (Fig. 46). The iodine, which col- 
lects in these condensers, is purified by washing and resubliming. 
(2) In another process the mother liquor from the Chili saltpeter is 
mixed with acid sodium sulphite (HNaSO 3 ), and the precipitated iodine 
is collected on coarse cloth, washed, dried, and then resublimed, as 
described above. 

Courtois, a French chemist, discovered iodine, in 1812, in an attempt 
to prepare potassium nitrate from seaweed. Davy and Gay-Lussac 
established its elementary nature and discovered many of its properties. 
The present name was given by Davy. 

Properties. Iodine is a dark grayish crystalline solid, 
resembling graphite in luster. It crystallizes in plates 
which have the specific gravity 4.95. It is volatile at the 

232 Descriptive Chemistry. 

ordinary temperature, and when gently heated the vapor 
which is formed has a beautiful violet color. This color 
suggested the name iodine (from the Greek word iodes, 
violetlike). The vapor is nearly nine times heavier than 
air, and has an odor resembling dilute chlorine, though less 
irritating. When the vapor is heated, its color changes 
from violet to deep blue, and the density decreases. Ex- 
periment indicates that at about 700 C. the molecules con- 
tain only two atoms, and as the temperature rises the 
molecules dissociate, until at a very high temperature the 
vapor consists entirely of atoms. Iodine stains the skin 
yellow, and turns cold starch solution blue. The presence 
of a minute trace of iodine may be thus detected, one part 
of iodine in over 400,000 parts of water producing the blue 
color. The exact nature of this blue compound is un- 
known. The presence of starch in many vegetable sub- 
stances can be shown by this delicate test. Iodine dissolves 
slightly in water, and freely in alcohol, chloroform, carbon 
disulphide, ether, and potassium iodide solution. The 
chloroform and carbon disulphide solutions are violet, but 
the others are brown, or even black. The chemical proper- 
ties of iodine resemble those of chlorine and bromine, but 
it is less active. Bromine and chlorine displace iodine 
from its compounds, chlorine and chlorine water being 
often used for this purpose. It combines directly with 
other elements and replaces some. Phosphorus bursts into 
a flame when mixed with iodine. 

Compounds of Iodine resemble the corresponding ones of chlorine 
and bromine. Hydriodic acid is much like hydrobromic and hydro- 
chloric acid, though unlike them in being a reducing agent. Iodides 
are salts of hydriodic acid, and like many salts they are prepared in 
various ways. In general behavior they are similar to bromides and 
chlorides. Potassium iodide (KI) is made and used like potassium 
bromide. lodates and periodates are known. 

Fluorine Bromine-^- Iodine. 233 

Miscellaneous. Iodine dissolved in alcohol or in potassium iodide 
solution is used as an application for the skin to prevent the spread of 
eruptions or to reduce swellings. Iodine is used to make medicinal 
preparations, especially iodoform (CHI 3 ), which is used as a dressing 
for wounds. Large quantities of iodine are used in making aniline 
dyes. Potassium iodide is made in large quantities, Germany alone 
exporting about 150 tons of it annually. Chili annually exports over 
300 tons and Norway over 160 tons of iodine and iodides. 


1. What elements constitute the halogen group ? Why are they 
so .called ? 

2. How does fluorine occur in nature ? Describe briefly the isola- 
tion of fluorine. When was it first performed? Summarize the chief 
properties of fluorine. 

3. How is hydrofluoric acid prepared? Give the equation for the 
reaction. What are its characteristic properties ? For what is it used ? 

4. How is glass etched? State the essential changes. 

5. What is the formula of hydrofluoric acid ? 

6. How does bromine occur in nature ? What are the sources of 
commercial bromine ? What general method is used to prepare this 
element ? Describe briefly the commercial methods. State the chief 
properties. For what is it used ? How does this element differ from 
all others previously studied ? 

7. Name several compounds of bromine. What is potassium 
bromide ? 

8. Give a brief account of the discovery of (a) bromine and 
(b) iodine. 

9. Discuss the occurrence of iodine in nature. How is iodine pre- 
pared (a} in the laboratory and (b) on a large sqale ? Summarize the 
properties of iodine. Describe the test for iodine. 

10. Name several compounds of iodine. Describe potassium iodide, 
n. Compare hydrochloric, hydrobromic, and hydriodic acids. 

12. What is the symbol of (a} fluorine, (b} chlorine, (c) bromine, 
(d) iodine ? What is the derivation of the name of each element ? 

13. Compare the physical properties of fluorine, chlorine, bromine, 
and iodine. 

14. What is " drug-store iodine " ? 

234 Descriptive Chemistry. 


1 . What is the percentage composition of (a) fluor spar (CaF 2 ) and 
() cryolite (Na 3 AlF 6 )? 

2. How much (a) calcium sulphate and () hydrofluoric acid are 
formed by heating 100 gm. of fluor spar with sulphuric acid ? 

3. Calculate the percentage composition of (a) potassium bromide 
(KBr), () potassium iodide (KI), (c} silver bromide (AgBr), and 
(</) iodoform (CHI 3 ). 

4. How much potassium iodide is needed to prepare 63.5 gm. of 
iodine ? 

5. How much potassium bromide is needed to prepare 10 gm. of 
bromine ? 


SULPHUR has been known for ages. The alchemists re- 
garded it as one of the primary forms of matter. The ele- 
ment and its compounds have always played an important 
part in the development of many industries. 

Occurrence and Formation. Sulphur, free and com- 
bined, is abundant and widely distributed. Free or native 
sulphur is found usually in volcanic regions. There are 
also beds associated with gypsum (calcium sulphate). It is 
believed that such deposits were formed by the reduction of 
the gypsum by microorganisms into limestone and sulphur. 

Combined sulphur is found in volcanic gases, in sub- 
stances of vegetable and animal origin, and as sulphides 
and sulphates. Several important metallic ores are native 
sulphides, e.g. lead sulphide (PbS), zinc sulphide (ZnS), and 
those of mercury, antimony, and copper. Probably some 
native sulphur has been formed by the decomposition of 
sulphides by heat. The most abundant sulphates are 
varieties of calcium sulphate (CaSO 4 ), barium sulphate 
(BaSO 4 ), and magnesium sulphate (MgSO 4 ). Volcanic 
gases often contain sulphur dioxide (SO 2 ) and hydrogen 
sulphide (H 2 S). The latter is also found in the water of 
sulphur springs. Doubtless some of the sulphur found in 
volcanic districts has been produced from these two gases. 
Their interaction may be represented thus 

SO 2 + 2 H 2 S = 3 S + 2H 2 O 

Sulphur Dioxide Hydrogen Sulphide Sulphur Water 



Descriptive Chemistry. 

Sulphur is also a component of onions, horse-radish, mus- 
tard, garlic, eggs, some petroleum and coal, and certain 
complex compounds of the body such as bile and saliva. 
It has been estimated that the body contains about 125 
gm. (0.27 Ib.) of combined sulphur. 

Source. Sicily furnishes most of the sulphur used in 
the world, the annual output being about 500,000 tons. 
Owing to the favorable geographical location, rich deposits, 
and cheap labor, the bulk of the supply will continue to 
come from this island. Some sulphur is obtained from 
Japan, Italy, Greece, and from the United States, especially 
in Nevada, Utah, Idaho, and Louisiana. 

Some of the sulphur of commerce is obtained by roasting iron pyrites, 
as in the manufacture of sulphuric acid. Small amounts are recovered 
from the calcium sulphide waste of the Leblanc soda process (see 
Sodium Carbonate), and from the residues of the iron oxide used to 
purify illuminating gas. . 

FIG. 47. Kiln for extracting sulphur from the crude ore. The calcarone is shown 
as a vertical section (right) and in operation (left). 

Extraction. For many years sulphur has been ex- 
tracted from the impure native sulphur in Sicily by a 
primitive process. The crude sulphur is brought to the 
surface by laborers, piled loosely in a heap, and covered 
with powdered or burnt ore or with earth. The heap is 
ignited at the bottom, and the heat produced by the com- 

Sulphur and its Compounds. 


bustion of some of the sulphur melts the rest, which runs 
out at the bottom (Fig. 47). 

This method is being discarded in the more prosperous localities, 
because it is wasteful and produces intolerable fumes. Coal instead of 
sulphur is being used as a fuel, and extraction by hot water under 
pressure is coming into general use. In some cases the sulphur is 
extracted by heating the crude sulphur with a hot solution of calcium 

Purification. Sulphur obtained from its ore requires 
purification. This is accomplished by the apparatus shown 
in Figure 48. The crude sulphur is melted in B, and flows 
into the iron cylinder, A. Here it is heated, and the vapors 

FIG. 48. Apparatus for purifying sulphur. 

pass into the large brick chamber, provided with a tap, C, 
from which the liquid sulphur may be withdrawn. If the 
distillation is conducted slowly, the sulphur vapor con- 
denses upon the cold walls of the chamber as a fine 

238 Descriptive Chemistry. 

powder, called flowers of sulphur, just as water vapor 
suddenly cooled below o C. turns to snow. As the 
operation continues the walls become hot, and the sulphur 
collects on the floor as a liquid which is drawn off into 
wooden molds. This is roll sulphur or brimstone. 

Properties, Ordinary sulphur is a yellow, brittle, crys- 
talline solid. It is insoluble in water, but most varieties 
dissolve in carbon disulphide, and to some extent in turpen- 
tine, chloroform, and benzene (C 6 H 6 ). Sulphur does not 
conduct heat. The warmth of the hand causes it to crackle 
and even break from the unequal expansion. 

The specific gravity of the solid is about 2. The specific gravity of 
the vapor varies with the temperature. At the lowest temperature at 
which sulphur can be vaporized, the molecule contains eight atoms 
(S 8 ), while at 900 C. and higher it contains two atoms (S 2 ). 

Heated to 1 14.5 C. sulphur melts to a thin, amber-colored 
liquid. As the temperature is raised, the liquid darkens 
and thickens, until at about 230 C. it is black and too thick 
to be poured from the vessel. Heated still higher, the 
color remains black but the mass becomes thin, and finally 
at about 448 C. the liquid boils and turns into a yellowish 
brown vapor. Sulphur ignites readily and burns with a 
pale blue flame, forming sulphur dioxide gas, SO 2 ; if 
burned in oxygen, a little sulphur trioxide, SO 3 , is also 
formed. Finely divided sulphur oxidizes in moist air, 
forming sulphuric acid, H 2 SO 4 . It also combines directly 
and readily with hydrogen, carbon, chlorine, and other ele- 
ments, especially metals. The compounds formed are 

The reaction between sulphur and metals is often attended by vivid 
combustion, though heat is necessary to start the chemical action. 
When a mixture of flowers of sulphur and powdered iron is heated, the 
mass begins to glow and soon becomes red-hot, the glow often spread- 

Sulphur and its Compounds. 239 

ing through the mass after removal from the flame. The product is 
iron sulphide, and the change is represented thus 

Fe + S FeS 

Iron Sulphur Iron Sulphide 

Heated copper glows when dropped into melted sulphur, while zinc 
dust and flowers of sulphur combine with almost explosive violence. 

Different Forms of Sulphur. Sulphur exists in at least 
three different forms, two crystallized and one amorphous. 
These modifications differ in specific gravity, solubility, 
and other properties. The crystallized forms belong to 
the orthorhombic and monoclinic systems (see Appendix, 
3). According to some authorities these different forms 
are allotropic modifications of sul- 
phur. Orthorhombic sulphur is the 
form deposited by crystallization 
from a solution of carbon disulphide 
(Fig. 49). Crystallized native sul- 
phur is orthorhombic. The mono- 
clinic crystals are deposited from 
molten sulphur. By melting sul- 
phur in a crucible and pouring off FIG. 49. Orthorhombic 
the excess of liquid as soon as crys- sulphur, 

tals shoot out from the walls near the surface, the interior of 
the crucible when cold will be found to be full of long, dark 
yellow, shining needles. They are monoclinic crystals of 
sulphur. After a few days they become dull and yellow, 
and crumble into minute crystals of the orthorhombic form. 

Amorphous sulphur is formed by pouring boiling sul- 
phur into water. It is a tough, plastic, rubberlike, amber- 
colored mass, insoluble in carbon disulphide. It is entirely 
different in color and texture from the crystallized varieties. 
In a short time it becomes hard, brittle, and yellow, like 
ordinary sulphur. 

240 Descriptive Chemistry. 

Other varieties of amorphous sulphur are known. They are white 
or whitish powders. One is made by boiling flowers of sulphur with 
milk of lime and adding hydrochloric acid to the decanted liquid. A 
fine sulphur powder is precipitated, which gives the liquid the appear- 
ance of milk, hence the name often applied to it, "milk of sulphur." 

Uses. Sulphur is used in making sulphuric acid and 
other sulphur compounds, gunpowder, fireworks, matches, 
in vulcanizing rubber, as a medicine and a constituent of 
some ointments, and as a germicide for Phylloxera an 
insect which destroys grapevines. 

Compounds of Sulphur. The important compounds of 
sulphur are hydrogen and other sulphides, sulphur dioxide 
and trioxide, the sulphites, sulphuric acid and the sulphates, 
and carbon disulphide. 

Hydrogen Sulphide, H 2 S, is a gaseous compound of 
sulphur and hydrogen, and is often called sulphuretted 
hydrogen. It occurs in some volcanic gases, and in the 
waters of sulphur springs. It is often found in the air, 
especially near sewers and cesspools, since it is one prod- 
uct of the decay of organic substances containing sul- 
phur. It is one of the impurities of illuminating gas, being 
formed by the union of the sulphur and hydrogen of the 

The gas is prepared in the laboratory by the interaction 
of dilute acids and metallic sulphides, usually hydrochloric 
acid and ferrous sulphide. When the acid is poured upon 
fragments of the sulphide, the gas is rapidly evolved with- 
out applying heat, and may be collected over water. The 
equation for the chemical change is 

FeS + 2HC1 = H 2 S + FeCl 2 

Iron Hydrochloric Hydrogen Iron 

Sulphide Acid Sulphide Chloride 

Sulphur and its Compounds. 241 

Hydrogen sulphide gas is colorless and has the odor of 
rotten eggs. It is poisonous. A little, if breathed, produces 
headache and nausea, and a large quantity renders one un- 
conscious. This gas is inflammable and burns with a bluish 
flame, forming water and sulphur dioxide, thus 

2H 2 S + 3O 2 = 2SO 2 + 2H 2 O 
Hydrogen Sulphide Oxygen Sulphur Dioxide Water 

If the supply of air is insufficient, combustion is incom- 
plete and sulphur is also formed. It is a powerful reduc- 
ing agent, and is often used as such in chemical analysis. 
Even sulphuric acid is reduced by it, thus 

H 2 S0 4 + H 2 S = S0 2 + S 4- 2 H 2 O 

Sulphuric Hydrogen Sulphur Sulphur Water 

Acid Sulphide Dioxide 

Hydrogen sulphide is soluble in water, one volume of water 
dissolving about three volumes of the gas at the ordinary 
temperature. The solution is called hydrogen sulphide 
water, and is often used instead of the gas. The solution 
reddens litmus and decomposes slowly, sulphur being de- 

A liter of dry hydrogen sulphide gas, under standard conditions, 
weighs 1.542 gm. When metals are heated in dry hydrogen sul- 
phide, metallic sulphides are formed and the volume of hydrogen liber- 
ated is the same as the original volume of gas. Since the hydrogen 
molecule is H 2 , there must be two atoms of hydrogen in the hydrogen 
sulphide molecule. Its vapor density is 17.15, hence the molecular 
weight is 34.3. Subtracting 2 for H 2 , the remainder 32.3 agrees well 
with the atomic weight of sulphur. Hence, there can be only one 
atom of sulphur in hydrogen sulphide, and formula must be H 2 S. 

Sulphides may be regarded as salts of the weak acid, 
hydrogen sulphide, though they are not always prepared 
directly from hydrogen sulphide. They may be produced 

242 Descriptive Chemistry. 

by the direct union of sulphur and metals, as in the case 
of iron and copper sulphides previously mentioned, or by 
exposing the metal to the moist gas. A more common 
way is to precipitate them by passing the gas into solutions 
of metallic compounds, or, sometimes, by adding. hydrogen 
sulphide water. Copper, tin, lead, and silver are rapidly 
tarnished by the gas. Silverware, on this account, turns 
brown or black, especially in houses heated by coal and 
lighted by coal gas, because hydrogen sulphide is one 
product of the combustion of coal and gas. The brown 
silver sulphide also coats silver spoons which are put into 
mustard or eggs. Lead compounds are blackened by this 
gas, owing to the formation of lead sulphide, thus 

PbO + H 2 S PbS + H 2 

Lead Oxide Hydrogen Sulphide Lead Sulphide Water 

For this reason houses painted with " white lead " paint 
often become dark, and, similarly, oil paintings are dis- 
colored. The blackening of a solution of a lead com- 
pound is the customary test for hydrogen sulphide. 

Many sulphides have a brilliant color. Arsenious sulphide is pale 
yellow, cadmium sulphide is golden yellow, manganese sulphide is flesh 
colored, zinc sulphide is white, antimony sulphide is orange red. They 
vary in solubility. The sulphides of lead, silver, copper, and some 
other metals are insoluble in dilute hydrochloric acid. The sulphides 
of iron, zinc, and some other metals are decomposed by dilute hydro- 
chloric acid, but are precipitated if ammonium hydroxide is present. 
Sulphides of certain metals dissolve in water. Hence by precipitating 
metals under different conditions, groups of metals may be separated 
and subjected to further tests. The color often affords a ready means 
of detecting each sulphide. Hydrogen sulphide is thus a serviceable 
reagent in the branch of chemistry called Qualitative Analysis. 

Sulphur Dioxide, SO 2 , is the common compound of sul- 
phur and oxygen. It occurs in the gases of volcanoes, 

Sulphur and its Compounds. 243 

and to a slight extent in the atmosphere, since it is the 
usual product of the combustion of sulphur and sulphur 

When sulphur burns in air (or oxygen), sulphur dioxide 
is formed, thus 

S + O 2 SO 2 

Sulphur Oxygen Sulphur Dioxide 

It is also formed by roasting iron disulphide (iron pyrites) 
in the air, thus 

2 FeS 2 + 1 1 O = 4 SO 2 + Fe 2 O 3 

Iron Disulphide Oxygen Sulphur Dioxide Iron Oxide 

The above reaction is utilized on a large scale in the com- 
mercial manufacture of sulphuric acid. 

Sulphur and carbon reduce sulphuric acid to sulphur dioxide, thus 

S + 2 H 2 SO 4 3 SO 2 +. 2 H 2 O 

Sulphur Sulphuric Acid Sulphur Dioxide 

C 4- 2H 2 SO 4 = 2SO 2 + CO 2 + 2H 2 O 
Carbon Carbon Dioxide 

Two methods of preparation are used in the laboratory. 

(1) If copper and concentrated sulphuric acid are heated, 
a series of complex changes results finally in the evolution 
of sulphur dioxide. The equation is usually written 

Cu + 2H 2 SO 4 = SO 2 + CuSO 4 + 2H 2 O 
Copper Sulphuric Acid Sulphur Dioxide Copper Sulphate 

(2) Dilute sulphuric (or hydrochloric) acid dropped upon 
a sulphite yields sulphur dioxide, thus 

Na 2 SO 3 + H 2 SO 4 = SO 2 + Na 2 SO 4 + H 2 O 

Sodium . Sulphuric Sulphur Sodium 

Sulphite Acid Dioxide Sulphate 

244 Descriptive Chemistry. 

This method is convenient for liberating a steady current 
of the gas. 

Sulphur dioxide gas has no color. Its odor is suffocating, 
being the well-known odor associated with burning sulphur 
matches. It will not burn in the air, nor will it support 
ordinary combustion. A burning taper or stick of wood 
is instantly extinguished by it, but finely divided metals, 
iron for example, burn in it. It is a heavy gas, the high 
density (2.2) allowing it to be readily collected by down- 
ward displacement. Low temperature and pressure change 
it into a transparent, colorless liquid, which boils at 8 C. 
and freezes at 76 C. into a transparent, icelike solid. 
It is very soluble in water. At the ordinary temperature 
one volume of water dissolves about forty volumes of gas, 
but loses it all by boiling. This solution is sour and red- 
dens blue litmus, and contains sulphurous acid. Moist 
sulphur dioxide bleaches vegetable coloring matters. A 
red or a purple flower loses color in it. Silk, hair, straw, 
wool, and other delicate substances, which would be injured 
by chlorine, are whitened by sulphur dioxide. In some 
cases the color returns when the bleached article is exposed 
to the air for some time, and usually such bleached objects 
become yellow with age. The coloring matter is not wholly 
destroyed, but probably unites with the sulphur dioxide to 
form a colorless compound, which slowly decomposes. 

Immense quantities of sulphur dioxide are used in the 
manufacture of sulphuric acid. The gas is also used to 
preserve meat and wines, to fumigate clothing and houses, 
in paper making, in tanning, in refining sugar, and in 
making acid sodium sulphite. Liquid sulphur dioxide is used 
in extracting glue and gelatine, and in various metallurgical 
processes. It absorbs heat during evaporation, and is used 
in some ice machines. 

Sulphur and its Compounds. 245 

A liter of sulphur dioxide under standard conditions weighs 2.868 gm. 

The Composition of Sulphur Dioxide is based on the following: 
The gas formed by burning sulphur in a measured volume of oxygen 
has the same volume as the oxygen itself. Hence there are as many 
molecules of sulphur dioxide as there were of oxygen ; that is, one 
molecule of sulphur dioxide contains one molecule (or two atoms) of 
oxygen. A molecule of oxygen weighs 32. But the molecular weight 
of sulphur dioxide found from its vapor density is about 64. Subtract- 
ing 32 (i.e. 2 x 1 6) from this, there remains about 32 for sulphur. The 
atomic weight of sulphur is 32.07, hence sulphur dioxide contains only 
one atom of sulphur, and its composition is expressed by the formula 
S0 2 . 

Sulphurous Acid and Sulphites. Sulphurous acid is formed when 
sulphur dioxide dissolves in water. Sulphur dioxide is, therefore, sul- 
phurous anhydride. The simplest equation expressing this fact is 

SO, + H 2 O H 2 S0 3 

Sulphur Dioxide Water Sulphurous Acid 

The acid has never been obtained free, resembling carbonic acid in this 
respect. It is unstable, and gradually forms sulphuric acid by combin- 
ing with oxygen from the air. The acid is dibasic, and forms two 
classes of salts, the sulphites. They are reducing agents, and yield 
sulphur dioxide when treated with acids. Acid sodium sulphite 
(HNaSO 3 ), often called bisulphite of soda, is the antichlor used to 
remove the excess of chlorine from bleached cotton cloth. It is also 
used in brewing, tanning, and in making starch, sugar, and paper. 
Acid calcium sulphite (CaH 2 (SO 3 ) 2 ), prepared by passing sulphur 
dioxide into milk of lime, is used in paper making. 

Sulphur Trioxide, SO 3 , is formed by the direct union of 
sulphur dioxide and oxygen, a little being produced when 
sulphur burns in air or in oxygen. The action is slow, but 
may be hastened by passing a mixture of sulphur dioxide 
and oxygen (or air) over hot platinum, or over asbestos 
coated with platinum. Other substances also hasten the 
change. It is a white, crystalline solid, which melts at 
15 C. and boils at 46 C. Another form, silklike in luster 
and appearance, is known. When exposed to moist air it 

246 Descriptive Chemistry. 

fumes strongly, forming sulphuric acid ; and when dropped 
into water it dissolves with a hissing sound and evolution 
of heat, thus 

SO 3 + H 2 O H 2 SO 4 

Sulphur Trioxide Water Sulphuric Acid 

The vapor density of sulphur trioxide shows that its molecular weight 
is about 80. Hence the formula (SO 3 ) harmonizes with the fact that 
two volumes of sulphur trioxide decompose by heat into two volumes of 
sulphur dioxide and one volume of oxygen. 

Sulphuric Acid, H 2 SO 4 , is found in the waters of a few 
rivers and mineral springs. It is manufactured in enormous 
quantities and used for many purposes. 

Sulphuric acid was doubtless known to the Arabian alchemists living 
in the tenth century. It was definitely mentioned by Basil Valentine 
in the fifteenth century, who describes its preparation by heating a mix- 
ture of iron sulphate (green vitriol) and sand. The product, an oily 
liquid, was called oil of vitriol, a name now often used. About 1740, 
the method of burning sulphur and oxidizing the product was introduced 
into England. 

The Manufacture of Sulphuric Acid, as usually con- 
ducted, is based upon the fact that the oxidation of sulphur 
dioxide in the presence of water forms sulphuric acid. 
The apparent equation for the chemical change is 

S0 2 4-0 4- H 2 = H 2 S0 4 

Sulphur Dioxide Oxygen Water Sulphuric Acid 

The oxidation is accomplished in the older and more com- 
mon method by oxides of nitrogen. T) * 

The general operation consists in passing sulphur diox- 
ide, air, steam, and oxides of nitrogen into large lead cham- 
bers. The oxides of nitrogen in the presence of steam 
change the sulphur dioxide into sulphuric acid, which col- 
lects on the walls and floors of the lead chambers. The 
oxides of nitrogen which lose part of their oxygen by this 

Sulphur and its Compounds. 


change are themselves reoxidized by the air into higher 
oxides, thus being fitted to oxidize more sulphur dioxide. 
The oxides of nitrogen act as carriers of oxygen, continu- 

ously giving oxygen to sulphur dioxide and taking it from 
the air. Theoretically, a small quantity of the oxides of 
nitrogen will change an infinite quantity of sulphur dioxide 

248 Descriptive Chemistry. 

into sulphuric acid, but in practice losses occur and oxides 
of nitrogen must be supplied. The main parts of a sul- 
phuric acid plant, together with the courses taken by the 
gases, are shown in Figure 50. 

Careful study shows that the chemical changes involved in this pro- 
cess of manufacturing sulphuric acid are complex and variable. Ac- 
cording to a reliable authority, the main continuous reactions may be 
represented thus 

2HNO 3 + 2SO 2 + H 2 = 2H 2 SO 4 + N 2 O 3 

2SO 2 -t- N 2 O 3 -f O 2 + H 2 O = 2 SO 2 (OH)(NO 2 ) 

Nitrosy 1-sul phuric 

2SO 2 (OH)(NO 2 ) + H 2 O= 2H,SO 4 + N,O 3 

or, 2SO 2 (OH)(NO 2 ) -f SO 2 + O + 2 H 2 O = 3 H 2 SO 4 + N,O 3 

The nitrogen trioxide (N 2 O 3 ) is the essential factor, though probably 
the change is really due to a mixture of nitric oxide (NO) and nitrogen 
peroxide (NO 2 ) . Under some conditions, nitric oxide plays a prominent 
part. It may be said in general that the ease with which the oxides of 
nitrogen pass into each other makes it highly probable that they are 
carriers of oxygen from the air to the sulphur dioxide. 

A Sulphuric Acid Plant consists of three main parts (a) the furnace 
for producing sulphur dioxide, () the lead chambers together with the 
Glover and Gay-Lussac towers for changing the sulphur dioxide into 
sulphuric acid, and (V) the concentrating apparatus. The manufacture 
is conducted somewhat as follows: (i) Sulphur or iron disulphide 
(FeS 2 ) is burned in a furnace constructed so that enough air passes 
over the burning mass to change the sulphur into sulphur dioxide, and 
to furnish the proper amount of oxygen for later changes. In some 
works the furnace is provided with " niter pots " containing a mixture 
of sodium nitrate and sulphuric acid ; the nitric acid vapors which are 
formed are one source of the oxides of nitrogen. (2) The mixture of 
sulphur dioxide, oxides of nitrogen, and air passes from the furnace into 
the bottom of the Glover tower. This is a tall tower filled with small 
stones over which flow two streams of sulphuric acid, one dilute and the 
other containing oxides of nitrogen (obtained from the Gay-Lussac 

Sulphur and its Compounds. 249 

tower). These acids not only cool the ascending gases, but are them- 
selves deprived of water and oxides of nitrogen. Hence, concen- 
trated acid flows out of the bottom of the Glover tower, while from the 
top sulphur dioxide, oxides of nitrogen, steam, and air pass on into the 
first lead chamber. Here nitric acid is often introduced, as well as 
steam. The main chemical changes occur in this and in the second 
chamber. A third chamber serves mainly to cool and dry the gases. 
These chambers are huge boxes often having a total capacity of 150,000 
cubic feet ; the walls and floors are of sheet lead supported on a wooden 
framework, lead being a metal which is only slightly attacked by the 
chamber acid. The remaining gases pass on into the bottom of the 
Gay-Lussac tower. This tower is filled with coke over which flows 
concentrated sulphuric acid (from the Glover tower), which absorbs the 
unused oxides of nitrogen. These oxides are liberated again in the 
Glover tower, hence there is little loss. At the end of the plant is a 
tall chimney, which serves as an exit for unused gases (such as nitro- 
gen) and also creates a draft strong enough to carry the gases through 
the chambers and tower. (3) The acid which is produced in the cham- 
bers and drawn off. from them at intervals contains about 67 per cent 
of the compound H 2 SO 4 . Ordinary commercial sulphuric acid which 
contains about 96 to 98 per cent is prepared from the chamber acid by 
evaporation, first in lead pans and finally in a platinum or an iron 

Another method of manufacturing sulphuric acid has 
recently been perfected, called the contact method. Sul- 
phur dioxide and air, carefully purified and properly cooled, 
are led through pipes containing plates covered with a 
contact mixture, which is chiefly finely divided platinum. 
The sulphur dioxide is oxidized to sulphur trioxide, thus 

SO 2 + O = SO 3 

Sulphur Dioxide Oxygen Sulphur Trioxide 

The sulphur trioxide is conducted into dilute sulphuric acid 
or water, thus producing a pure acid of any desired strength. 
The process is continuous if the gases from the pyrites 
burners are completely freed from arsenic compounds, sul- 
phur dust, and other impurities. 

250 Descriptive Chemistry. 

In the above process the platinum is not changed, nor does it cause 
the sulphur dioxide to unite with the oxygen. It facilitates the chemi- 
cal action between the gases somewhat as oil assists the movement of 
machinery. This kind of chemical action is called catalysis or cata- 
lytic action. The substance which hastens or retards a chemical 
reaction, but appears unchanged at the end of the process is called a 
catalyzer. In many cases of catalytic action it has been found that the 
catalyzer probably participates in the chemical action, though its exact 
share is not always clearly understood. 

Properties of Sulphuric Acid. Sulphuric acid is an 
oily liquid, colorless when pure, but usually brown from 
the presence of charred organic matter, such as dust and 
straw. The commercial acid has the specific gravity 1.83. 
When sulphuric acid is mixed with water, considerable 
heat is evolved. The acid should always be poured into 
the water, otherwise the intense heat may crack the vessel 
or spatter the hot acid. The volume of dilute acid pro- 
duced is smaller than the sum of the volumes of water and 
concentrated acid. The tendency to absorb water is shown 
in many ways. The concentrated acid absorbs moisture 
from the air and from gases passed through it. It is often 
used in the laboratory to dry gases, since it is not volatile 
at the ordinary temperature. Wood, paper, sugar, starch, 
cotton cloth, and many organic substances are blackened by 
sulphuric acid. Such compounds contain hydrogen and 
oxygen in the proportion to form water ; these two ele- 
ments are abstracted and carbon alone remains. Similarly, 
sulphuric acid withdraws water from the flesh, making 
painful wounds. 

Sulphuric acid is reduced by hydrogen sulphide, hydrobromic and 
hydriodic acids, carbon, and sulphur ; it combines with ammonia to form 
ammonium sulphate (NH 4 ) 2 SO 4 ; and is decomposed by all metals ex- 
cept platinum and gold, liberating hydrogen, sulphur dioxide, or hydro- 
gen sulphide. 

Sulphur and its Compounds. 251 

Uses of Sulphuric Acid. Sulphuric acid is one of the 
most important substances. Directly or indirectly it is 
used in hundreds of industries upon which the comfort, 
prosperity, and progress of mankind depend. It is used 
in the manufacture of all other mineral acids and many 
organic acids. It is essential in one process for the manu- 
facture of sodium carbonate, from which in turn are made 
soap and glass. Enormous quantities are consumed in 
making artificial fertilizers, alum, nitroglycerine, glucose, 
phosphorus, dyestuffs, and in various parts of such funda- 
mental industries as dyeing, bleaching, electroplating, 
refining, and metallurgy. 

Sulphates. JSulphuric acid is dibasic and forms two 
classes of salts, the normal sulphates, such as Na 2 SO 4 , 
and the acid sulphates, such as HNaSO 4 . The normal 
sulphates are stable salts ; the acid salts lose water when 
heated. Most sulphates are soluble in water, only the sul- 
phates of barium, strontium, and lead being insoluble, 
while calcium sulphate is slightly soluble. Important 
sulphates are calcium sulphate (gypsum CaSO 4 .2 H 2 O), 
barium sulphate (heavy spar, BaSO 4 ), zinc sulphate (white 
vitriol, ZnSO 4 ), copper sulphate (blue vitriol or blue stone, 
CuSO 4 ), iron sulphate (green vitriol, copperas, ferrous sul- 
phate, FeSO 4 ), sodium sulphate (Glauber's salt, Na 2 SO 4 ), 
and magnesium sulphate (Epsom salts, MgSO 4 ). Sul- 
phates are widely used in medicine and in many industries. 

The test for sulphuric acid or a soluble sulphate is the formation 
of the white, insoluble barium sulphate upon the addition of barium 
chloride solution. An insoluble sulphate fused on charcoal is reduced 
to a sulphide, which blackens a moist silver coin. 

Fuming Sulphuric Acid, H 2 S 2 O 7 , is made by adding sulphur trioxide 
to sulphuric acid, or by heating moist ferrous sulphate. This is the 
acid called sulphuric acid by the alchemists. It is sometimes called 

252 Descriptive Chemistry. 

Nordhausen sulphuric acid. It is a thick, brown liquid, which fumes 
strongly in the air, owing to the escape of oxides of sulphur. It is used 
in gas analysis to absorb ethylene and other illuminants, and in dyeing 
to dissolve indigo. If the fuming acid is cooled to o C, crystals sepa- 
rate ; they are called pyrosulphuric acid. 

Sodium Thiosulphate, Na 2 S 2 O 3 , is a salt of an unstable acid. It is 
sometimes incorrectly called sodium hyposulphite, or simply " hypo." 
It is a white, crystallized solid, very soluble in water. The solution, 
used in excess, dissolves the halogen compounds of silver ; hence its 
extensive use in photography (see Photography). It also finds some 
use as an antichlor, and in chemical analysis for determining the amount 
of free iodine in a solution. 

Carbon Disulphide, CS 2 , when pure, is a clear, colorless liquid, with 
an agreeable odor. The commercial substance is yellow and has an 
offensive odor. It is poisonous. It is volatile and extremely inflam- 
mable, the equation for its combustion being - 

CS 2 + 3 2 = C0 2 + 2 SO, 

Carbon Disulphide Oxygen Carbon Dioxide Sulphur Dioxide 

This liquid is insoluble in water. It dissolves rubber, gums, fats, resins, 
iodine, camphor, and some forms of sulphur. It is a highly refracting 
liquid, and hollow glass prisms filled with it are used to decompose 
light. As a solvent it is used to dissolve pure rubber in the manufac- 
ture of rubber cement. It is also used to kill insects on both living and 
dried plants (e,g. in museums), and to exterminate burrowing animals, 
such as moles and woodchucks. Many oils, waxes, and greases are ex- 
tracted by carbon disulphide. It is also used to manufacture compounds 
of sulphur and of carbon. 

Until recently carbon disulphide was manufactured by passing sul- 
phur vapor over red-hot coke or charcoal in iron or earthenware retorts ; 
the product required laborious purification. It is now manufactured by 
an electrothermal process. Several groups of carbon electrodes are set 
into the base of a furnace, coke is packed loosely around them, and the 
body of the furnace is filled with charcoal. Sulphur is introduced at 
suitable points, and when the current passes the sulphur melts, vapor- 
izes, and unites with the heated carbon above the electrodes. 

Selenium and Tellurium are rare elements which form compounds 
analogous to the principal compounds of sulphur. These three with 
oxygen form a natural group, their physical properties varying gradually 
with increasing atomic weight. 

Sulphur and its Compounds. 253 


1. What is the symbol and atomic weight of sulphur? 

2. Where is free sulphur found? Discuss its formation. In what 
forms is combined sulphur found? Name five native compounds of 
sulphur. What animal and vegetable compounds contain sulphur? 

3. Give a brief account of the sulphur industry in Sicily. How is 
sulphur purified? 

4. What is (a} flowers of sulphur, (<) brimstone, (^) roll sulphur, 
(a) milk of sulphur? 

5. Summarize the properties of sulphur, especially its action when 

6. Describe the different forms of sulphur. 

7. For what is sulphur used? 

8. What is hydrogen sulphide? Where is it found? Describe its 

9. Summarize the properties of hydrogen sulphide. State the 
equation for its combustion. What is its action upon sulphuric acid?^ 
What is hydrogen sulphide water? 

10. Why is H 2 S the formula of hydrogen sulpMde2__ ^ 
n. What are sulphides? How are they formed? Name and de- 
scribe five. Why does silverware often blacken? What use is made 
of sulphides in qualitative analysis? 

12. What is sulphur dioxide? How is it formed? State one equa- 
tion for its formation. Describe its preparation. For what is it used? 

13. Summarize the properties of sulphur dioxide. 

14. Why is SO 2 the formula of sulphur dioxide? 

15. What is the volumetric equation for the formation of sulphur 
dioxide from sulphur and oxygen? How many liters of oxygen are 
needed to form 5 1. of sulphur dioxide? 

1 6. Discuss sulphurous acid and sulphites. , 

17. What is sulphur trioxide? How is it prepared? State its chief 
properties. What is its formula? Why? 

1 8. Give a brief historical account of sulphuric acid. Why is it often 
called oil of vitriol ? What is (a) chamber acid, (b) Nordhausen acid, 
(c} fuming sulphuric acid, (d) pyrosulphuric acid? 

19. Upon what fact is the manufacture of sulphuric acid based? In 
what two general ways is the operation accomplished ? 

20. Describe the older method of manufacturing sulphuric acid. 

254 Descriptive Chemistry. 

21. Describe the contact method of manufacturing sulphuric acid. 

22. Define (a) catalysis and (b} catalyzer. 

23. Summarize the properties of sulphuric acid. 

24. Enumerate the important uses of sulphuric acid. 

25. Define and illustrate (a) sulphate, () normal sulphate, (<:) acid 

26. What is (a) gypsum, (b) white vitriol, (c) green vitriol, (d) blue 
vitriol, (/) Glauber's salt, (_/") kieserite? 

27. Describe the test for (a) sulphuric acid, (b} sulphurous acid, 
(c) a soluble sulphate, (d) an insoluble sulphate, (tf) a sulphite. 

28. State (a) the properties, and ($) the uses of sodium thiosulphate. 
What is its common name? 

29. State (a} the properties, and (b) the uses of carbon disulphide. 
How is it manufactured? 


1. Calculate the percentage composition of (a) barium sulphate 
(BaSO 4 ), () zinc sulphate (ZnSO 4 ), (c) sodium sulphate (Na,SO 4 ). 

2. Calculate the percentage composition of (a) galena (PbS), (b) 
zinc blende (ZnS), (c) iron pyrites (FeS 2 ), (d) ferrous sulphide (FeS). 

3. What weight and what volume of hydrogen can be obtained from 
102 gm. of hydrogen sulphide ? 

4. What is the weight of a stick of brimstone 10 cm. long and 4 
cm. in diameter ? 

5. How many grams of ferrous sulphide are needed to prepare a liter 
of hydrogen sulphide gas ? 

6. Sulphuric acid is i .8 times heavier than water. How many grams 
of acid will a liter flask hold ? 

7. Calculate the weight of oxygen necessary to burn (to sulphur di- 
oxide) 731 gm. of sulphur containing 15 per cent of impurities. 

8. A lump of sulphur weighing 32 gm. is burned in air. Calculate 
(a) the weight of oxygen required, and () the weight of sulphur di- 
oxide formed. 

9. How many liters of oxygen are needed (a) to form 10 1. of 
sulphur dioxide by burning sulphur in air, and () to change 10 1. of 
sulphur dioxide to sulphur trioxide ? 


Occurrence of Silicon. Silicon does not occur free in 
nature, being found almost exclusively as silicon dioxide 
(SiO 2 ) or as silicates. These compounds are so abundant 
and widely distributed that approximately one fourth of the 
earth's crust is silicon. Sand and the different varieties of 
quartz are silicon dioxide. Most rocks are silicates. 

Silicon itself is a rare element. It is obtained with difficulty by 
heating silicon dioxide with carbon, aluminium, or magnesium in the 
electric furnace, or by heating silicon chloride with sodium. 

Like carbon, silicon has three allotropic forms, a brown amorphous 
powder, a dark grayish mass like graphite, and steel-colored crystals. 
Amorphous silicon may be changed into the other forms. They have 
different properties. 

The name " silicon " comes from the Latin word silex, silicis, flint. 

Silicon Dioxide or Silica, SiO 2 , is the most common com- 
pound of silicon. Sand, gravel, sandstone, and quartzite 
are almost wholly silica. It is an essential ingredient of 
many rocks, as granite and gneiss. Quartz is silicon di- 
oxide. It has many varieties, which differ in color and 
structure, due to minute impurities or to the mode of 
formation. Among the crystalline varieties are the clear, 
colorless rock crystal, the purple amethyst, and the rose, 
yellow, glassy, milky, and smoky forms. Varieties imper- 
fectly crystalline or amorphous are the waxlike chalcedony, 
the various forms of agate having different colored layers, 
the reddish brown carnelian, the black and white onyx, the 

2 55 

Descriptive Chemistry. 

red or brown jasper, the dull brown or black flint, and the 
brittle chert. Opal is hydrated silica (SiO 2 nH 2 O). Petri- 
fied or silicified wood is largely some variety of quartz 
which has replaced the woody fiber. There is a " petrified 
forest " in Arizona. Infusorial or diatomaceous earth is a 
variety of silica consisting of the shells of minute organisms 
called diatoms (\g. 51). Quartz is often found as crystals 
which consist usually of a six-sided prism with a six-sided 

pyramid at one or both 
ends, but the crystals 
are sometimes complex 
(Fig. 52). 

Quartz crystals and 

FiG. 51. Earth from Richmond, Va., con- 
taining diatoms. 

FIG. 52. Quartz crystals. 

varieties like them are hard enough to scratch glass. 
They are insoluble in water and acids, except hydro- 
fluoric acid, but are soluble in melted hydroxides and 
carbonates of sodium and potassium. Quartz is infusible, 
except in the oxyhydrogen flame. If fused with certain 
precautions, the molten mass can be drawn out into elastic 
threads, which are used to suspend delicate parts of elec- 
trical instruments. 

Sandstone and quartzite are used as building stones, and 
hard sandstone is made into grindstones and whetstones. 
Sand is used in making sandpaper, glass, porcelain, and 

Silicon and Boron. 257 

mortar. Glass is roughened and cut by blowing or " blast- 
ing " fine sand against it. Many of the varieties of quartz 
are cut and polished into ornaments and gems, e.g. amethyst, 
opal, and agate. Rock crystal is used as the " diamond " 
in cheap jewelry, and is cut into lenses for eyeglasses and 
optical instruments. Petrified wood is cut and polished into 
table tops, mantelpieces, and fireplaces. Infusorial earth is 
used to polish silver, "electro-silicon" being the commercial 
name of one kind, and in making cement, " soluble glass," 
dynamite, and refractory brick. Over 1300 tons are 
annually used in the United States. 

Silica and Plants. Ashes of many plants contain silica, showing 
that some compound of silicon is assimilated by the plant from the soil 
probably silicic acid or a soluble silicate (see below). The ashes of 
rye and wheat straws and of potato stems contain from 40 to 70 
per cent of silica. Plants like horsetail, sword grass, and bamboo are 
rich in silica. The silica is probably not a plant food in the strict sense, 
but gives firmness to the tall stalks, especially to their joints, and pro- 
duces the tough exterior coating, as on the bamboo. The quills of 
feathers and the spikes of sponges are tough and rigid from the silica 
they contain. 

Silicon Tetrafluoride (SiF 4 ) is formed by the interaction of silicon 
dioxide and hydrofluoric acid, as described under etching (see Etching). 

Silicic Acid and Silicates. When silicon dioxide is 
fused with sodium or potassium carbonates, the correspond- 
ing silicate is formed thus 

SiO 2 + K 2 CO 3 = K 2 SiO 3 + CO 2 

Silicon Potassium Potassium Carbon 

Dioxide Carbonate Silicate Dioxide 

Potassium and sodium silicates dissolve in water, and when 
hydrochloric acid is added, the gelatinous precipitate 
formed is a silicic acid having the formula H 2 SiO 3 (proba- 

258 Descriptive Chemistry. 

bly). This acid is decomposed, by heating, into silicon 
dioxide and water, thus 

H 2 SiO 3 = SiO 2 + H 2 O 

Silicic Acid Silicon Dioxide Water 

There are many complex silicic acids. Silicates are salts 
of silicic acids, though they are often so complex that no 
actual corresponding acid is known. Silicates make up 
a large part of the earth's crust, silicates of aluminium, 
iron, calcium, potassium, sodium, and magnesium being the 
most abundant. Many common rocks and minerals are 
silicates, e.g. feldspar, mica, mica schist, hornblende, clay, 
slate, beryl, garnet, serpentine, and talc. 

Sodium and potassium silicates are the only ones soluble 
in water, and the thick, sirupy solution is often called 
"water glass " or soluble silica. It is used in making yel- 
low soaps, cements, and artificial stone, to fix colors in 
frescoing and calico printing, and to render cloth, wood, 
and paper fireproof. 

Some forms of silica dissolve in a hot solution of sodium carbonate. 
Hence, many hot springs, as in the Yellowstone Park, contain silica in 
solution (as an alkaline silicate), and when the water comes to the sur- 
face and cools, silica is deposited around the spring in beautiful forms 
called geyserite or siliceous sinter. Probably the formation of petri- 
fied wood is due to the deposition of silica from such a solution. 

Silicides are compounds of silicon and other elements. Carborun- 
dum, carbon silicide (or silicon carbide, CSi), has been mentioned (see 
Carborundum, Chapter X). Silicides of iron, chromium, and copper 
(Fe 2 Si, Cr 2 Si, and Cu 2 Si) are also commercially important. 

Glass is a mixture of silicates, one of which is always 
a silicate of potassium or sodium. Window glass is a sili- 
cate of sodium and calcium, and Bohemian glass is a 
silicate of potassium and calcium. In flint glass, calcium 
is replaced by lead. 

Silicon and Boron. 


Glass is not made by mixing silicates, but by melting 
together sand, an alkali, and a calcium or a lead compound. 
The alkali may be sodium carbonate (Na 2 CO 3 ), or potas- 
sium carbonate (K 2 CO 3 ), or a mixture of these; sodium 
sulphate is often used. The calcium compound used is cal- 
cium carbonate (CaCO 3 ) in the form of chalk or limestone. 
The lead compound used is litharge (PbO) or red lead 
(Pb 3 O 4 ). Small quantities of other substances are also 
used, e.g. broken glass to help lower the melting point of 
the mixture, oxide of arsenic (As 2 O 3 ), potassium nitrate 
(KNO 3 ), or manganese dioxide (MnO 2 ) to remove the 
greenish color caused by iron compounds, metallic oxides 
or other substances to -produce colored glass, and numerous 
ingredients, such as calcium fluoride or calcium phosphate, 
to make special kinds of glass. 

The process consists in heating the proper mixture in a 
fire-clay pot to a high temperature. During the melting, 
gases escape, and the impurities, which rise to the surface 
as a scum, are removed. The molten mass is alldwed to 
cool until it becomes pasty. In this condition it may be 
blown, welded, cut, drawn, or molded into almost any 
desired shape. 

The mixture used varies with the kind of glass to be made. A typi- 
cal mixture for table and bottle glass, used in a large works, is 

Sand 1550 Ib. 

Sodium carbonate . . . . .' 550 Ib. 

Lime 200 Ib. 

Sodium nitrate ....... 100 Ib. 

Total charge 2400 Ib. 

Window Glass is made by blowing a lump of glass into a hollow 
globe and then into a cylinder ; this on being opened at both ends and 
cut lengthwise spreads open flat. Plate glass, which has about the 

260 Descriptive Chemistry. 

same composition as window glass, is made by pouring the molten 
glass upon a large table, rolling it with a hot iron roller, and subse- 
quently grinding and polishing it. Plate glass is used for large win- 
dows and for mirrors, but considerable rough plate is used for skylights 
and floors. Crown glass is a good quality of window glass. It has a 
brilliant surface. Limited quantities are used as " bull's eyes " in deco- 
rative windows. Bohemian glass is the hard glass of which much 
chemical apparatus is made. Flint glass is a silicate of potassium and 
lead ; it is a lustrous, soft glass, largely used in making lamp chimneys 
and globes. Pure flint glass is often called strass or paste, and on ac- 
count of its luster and brilliancy it is made into artificial gems. Lenses 
for telescopes and other optical instruments usually consist of both 
crown and flint glass. Cut glass is flint glass. The object is first 
molded or blown into the general shape, the design is then cut into 
the soft glass by a wheel, and the finished object is polished by a 
wooden wheel smeared with rouge (oxide of iron) or putty. 

Many objects, such as tumblers and small dishes, are now made by 
pressing the soft glass with a die or by blowing it into a mold. Fruit 
jars, bottles, and lamp chimneys are blown by machinery. Many other 
improvements have increased the output and improved the quality of 

All glassware must be cooled slowly to prevent the glass from being 
brittle. This operation is called annealing, and is accomplished by 
passing the objects slowly through a furnace in which the temperature 
is gradually lowered. 

Glass is colored by adding different substances which dissolve in the 
molten mass. Iron and chromium compounds make it green, the green 
color of many bottles and fruit jars being due to the iron in the cheap 
materials used ; copper and cobalt compounds produce different shades 
of blue ; manganese dioxide gives a pink or a violet, and a mixture of 
manganese dioxide and iron oxide gives an orange color ; yellow is pro- 
duced by charcoal, sulphur, or silver; certain copper compounds or 
gold give a ruby color ; translucent or white glass is made by adding 
fluor spar or cryolite ; smoked glass contains nickel ; iridescent glass 
is made by exposing it to the vapors of hydrochloric acid or of tin 
chloride (SnCl 4 ). 

The United States produces yearly about 50,000,000 dollars' worth 
of glass. The industry is carried on in about twenty-five states, Penn- 
sylvania producing two fifths of the total output. 

Silicon and Boron. 261 


Occurrence. Boron is never found free, but the com- 
pounds, borax (Na 2 B 4 O 7 ) and boric acid (H 3 BO 3 ), are 

Boron itself is an uncommon element. It is prepared by heating 
the oxide (B 2 O 3 ) with magnesium, aluminium, sodium, or potassium. 
It is greenish brown amorphous powder, without taste or odor. It 
burns when heated in air, forming the oxide (B 2 O 3 ). It also unites 
with the halogens, sulphur, and nitrogen. It forms many borides, one 
of which, carbon boride (CB,.), is said to be harder than diamond. 

Boric Acid, H 3 BO 3 , is contained in the waters and steam 
of certain volcanic regions, notably Tuscany. Large 
basins or tanks are built around these steam jets, and 
are arranged so that the water flows at intervals from one 
reservoir into the next lower, constantly becoming charged 
with more boric acid, as the steam condenses. The final 
solution is evaporated by aid of the heat from the steam 
jets, and the crude boric acid ;which settles out is purified 
by recrystallization. This compound is sometimes called 
boracic acid. 

Considerable boric acid is also made in California from borax, and 
in Germany from the boracite found at Stassfurt. 

Boric acid crystallizes in lustrous, white flakes, which feel greasy. 
It dissolves slightly in cold water, readily in hot water, and in alcohol. 
When the alcoholic solution is burned, a boron compound colors the 
vapor green. This is the test for boron compounds. 

Boric acid is used in making borax, in the manufacture of enamels 
and glazes for pottery, as an antiseptic in medicine and surgery, and for 
preserving meat, fish, milk, butter, beer, and wine. 

Borax, Na 2 B 4 O 7 . ioH 2 O, occurs in large quantities in 
California, and an impure borax called tinkal comes from 
Tibet. Much of the commercial borax is made from 


Descriptive Chemistry. 

boric acid or from native calcium borate (colemanite, 
Ca 2 B 6 O n . 5 H 2 O) by boiling with sodium carbonate and 
separating the borax by crystallization. 

Borax is a white crystallized solid, having ten or five 
molecules of water of crystallization. It effloresces in the 
air. When heated, ordinary borax melts, then swells up 
into a white porous mass, which finally becomes a glassy 
solid. This glassy borax dissolves metallic substances, 
especially oxides. If the borax is melted on the end of 
a looped platinum wire, the transparent globule is called a 
borax bead. These beads differ in color under .different 
circumstances, and the oxides of metals cause the beads to 
assume colors which are characteristic of the metals, as 
may be seen by the following table : 









Chromium . 

Reddish yellow 

Yellowish green 



Cobalt . . 





Copper . . 


Greenish Blue 



Manganese . 





The bead test is often used in chemistry to confirm other 
observations or to suggest further examination. 

Borax is used in the manufacture of enamels and glazes, 
and in the formation of the " paste " for artificial gems. 
Immense quantities are used for preserving canned meat 
and fish. It is a cleansing agent, and large quantities are 
consumed in laundries as well as in the manufacture of 

Silicon and Boron. 263 

soaps, particularly those intended for use in hard water 
(see Soap). Its power to dissolve oxides adapts it for use 
in soldering metals. Solder adheres only to clean metals, 
so a little borax is used to dissolve the film of oxide on 
the surfaces to be joined. It is likewise used in welding 
metals and as a flux in their preparation. Considerable 
quantities are used as a mordant in calico printing and in 
dyeing. It is an ingredient of ointments, lotions, and 
powders, which are designed to relieve hoarseness or 
skin eruption. 


1. What is the symbol and atomic weight of (a) silicon, and 
(fr) boron ? 

2. How is silicon found in nature ? What proportion of the earth's 
crust is combined silicon ? 

3. Name several common forms of silicon dioxide. Describe the 
different varieties of quartz. 

4. What is (a) petrified wood, () opal, (c) diatomaceous earth, 
(d) " electro-silicon " ? 

5. Summarize the properties of quartz. How can it be readily 
distinguished from other minerals and rocks ? 

6. State the uses of the different forms of silicon dioxide. 

7. Discuss the relation of silicon dioxide to plants. 

8. Review with special reference to silicon compounds (a) car- 
borundum, and () etching glass. 

9. Describe the formation and state the properties of ordinary 
silicic acid. Name several common silicates. What metals are com- 
ponents of silicates ? 

10. Describe the formation, state the uses, and enumerate the prop- 
erties of "water glass." 

11. What is glass ? How is it made ? Name the components of 
the different kinds. 

12. What is (a) window glass, () plate glass, (c) Bohemian glass, 
(d) flint glass, and (e) cut glass ? 

13. How is glass (#) annealed, and ($) colored ? 

14. How is boron found in nature ? What is the formula of (#) 
borax, and () boric acid ? 

264 ' Descriptive Chemistry. 

15. Where is boric acid found ? How is it manufactured ? State 
its properties and uses. 

1 6. Where is borax found ? How is it prepared for commerce ? 
State its properties and uses. 

17. Describe the borax bead. State and illustrate its use. 


1 . Calculate the percentage composition, of (#) willemite (Zn 2 SiO 4 ), 
(b) steatite (Mg s Si 4 O H . H 2 O), (V) quartz (SiO 2 ). 

2. What per cent of borax (Na 2 B 4 O 7 . 10 H 2 O) is boron ? 


PHOSPHORUS, arsenic, antimony, and bismuth, together 
with nitrogen, form a natural group of elements. 


Occurrence. Free phosphorus is not found in nature, 
but phosphates are numerous and abundant. The most 
common are phosphorite (impure Ca 3 (PO 4 ) 2 ) and apatite 
(3 Ca 3 (PO 4 ) 2 .CaCl 2 or CaF 2 ). About o.i per cent of the 
earth's crust is phosphorus. Calcium phosphate is pres- 
ent in all fertile soils, being a product of decayed rocks. 
Plants and animals contain phosphorus compounds as 
essential constituents of the brain, nerves, and bones. 

Phosphorus was discovered in 1669 by Brand, who obtained it by 
heating a certain 'kind of animal matter. Scheele, in 1771, extracted it 
from bones. 

Preparation. Phosphorus is too dangerous a substance to prepare 
in the laboratory, (i) It is manufactured from bone ash or from native 
phosphates. The finely ground material is mixed in large vats with 
enough sulphuric acid to produce the following change: 

Ca,(P0 4 ) 2 + 3H,S0 4 = 2H 3 PO 4 + 3 CaSO 4 

Calcium Sulphuric Acid Phosphoric Acid Calcium 

Phosphate (Ortho-) Sulphate 

The insoluble calcium sulphate is removed by filtering the mixture 
through cinders. The phosphoric acid solution is concentrated, mixed 
with sawdust, coke, or charcoal, and dried, being changed thereby 
according to the equation 

H,PO 4 . HPO 3 + H 2 O 

Phosphoric Acid (Ortho-) Phosphoric Acid (Meta-) 



Descriptive Chemistry. 

The dried mass is heated to a high temperature in clay retorts arranged 
in tiers (Fig. 53), the change thus produced being substantially 

2H 2 + 12 CO 
Hydrogen Carbon 

4HPO 3 + 12 C = P 4 + 

Phosphoric Acid Carbon Phosphorus 

The phosphorus distils as a vapor through a pipe into a trough of water, 
where it collects as a heavy liquid. (2) Phosphorus is also manufactured 
in the electric furnace. A mixture of a phosphate, carbon, and sand is 
fed into a furnace provided with an outlet pipe through which the phos- 
phorus vapor passes into a condenser. The residue is drawn off as a 
slag at the bottom. The equation for the chemical change is 

2Ca 3 (P0 4 ) 2 


6SiO 2 + 10 C = P 4 + 10 CO + 3CaSiO 3 
Sand Carbon Phosphorus Carbon Calcium 

Monoxide Silicate 

Either method gives a black product, which is purified by redistil- 
lation in an iron retort, or by oxidation under water with sulphuric 

acid and potassium dichromate ; 
finally it is pressed through can- 
vas bags and molded into sticks. 

Properties. Phosphor- 
us has three allotropic 
modifications, yellow or 
ordinary, red or amorphous, 
and black or crystalline. 
Ordinary phosphorus is 
a yellowish, translucent 
solid. The color deepens 
by exposure to light. At 
ordinary temperatures 
phosphorus is like wax, 
but at low temperatures it 
is brittle. Under water it 
melts at 44 C. Exposed 

FIG. 53. Apparatus for the manufacture 

of phosphorus. to the air it immediately 

Phosphorus, Arsenic, Antimony, Bismuth. 267 

gives off white fumes, and at 34 C. takes fire and burns 
with a brilliant flame, the main product being phosphorus 
pentoxide (P 2 O 5 ). In moist air it glows, as may be easily 
seen by rubbing the head of a match in a dark room. 
This property gave the element its name (from the Greek 
word phosphoros, light bringer). The ease with which it 
ignites makes phosphorus dangerous to handle. Burns 
from it are severe and hard to heal. It is very poisonous, 
and the workmen in phosphorus factories are subject to 
a dreadful disease, which rots the bones. A fatal dose 
is about o. 1 5 gm. Phosphorus is kept beneath water, and 
should never be handled or cut unless so covered. It is 
nearly insoluble in water, but dissolves in carbon disulphide 
and slightly in sodium hydroxide solution. .Yellow phos- 
phorus has a faint odor, which may be easily detected by 
smelling a match head. Red phosphorus is made by 
heating ordinary phosphorus to 25O-3OO C. in a closed 
vessel. Any unchanged yellow phosphorus is extracted 
with sodium hydroxide solution. The red phosphorus is 
usually a reddish brown powder, though sometimes it is 
a brittle mass. It is opaque and odorless, does not give 
light, nor can it be easily ignited. It is poisonous, and 
does not dissolve in carbon disulphide. Its specific gravity 
is 2.25, that of the yellow form being 1.836. It can be 
handled without danger. Heated to about 260 C. in an 
atmosphere of nitrogen or carbon dioxide, it changes into 
ordinary phosphorus. 

Black Phosphorus is formed by dissolving red phosphorus in melted 
lead, and allowing crystals to separate. Its specific gravity is 2.34. 

The vapor density of phosphorus is such that its molecule must 
contain four atoms, hence its molecular formula is P 4 . 

Certain rat and bug poisons contain ordinary phosphorus, but most 
of the phosphorus of commerce is consumed in the manufacture of 
matches (see below). 

268 Descriptive Chemistry. 

Oxides of Phosphorus. The two important oxides are phosphorus 
or trioxide (P 2 O 3 or P 4 O 6 ) and phosphoric or pentoxide (P 2 O 5 ). Phos- 
phorous oxide is a white solid formed by the slow oxidation of phos- 
phorus or by burning phosphorus in a limited supply of air. It has the 
odor of phosphorus and is poisonous. Warmed in the air, it changes 
into the pentoxide. It unites with water to form phosphorous acid, 
thus- p^ + 3H2 o = 2 H,P0 3 

Phosphorous Oxide Phosphorous Acid 

Phosphoric Oxide (P 2 O 5 ) is the white, snowlike solid formed by 
burning phosphorus in an abundant supply of air. It is very deli- 
quescent, quickly withdrawing moisture from the air and combining vig- 
orously with water with a hissing noise. It resembles sulphur trioxide 
in its power to char wood and paper by withdrawing from them the 
elements of water. It is often used in the laboratory to dry gases. 

Acids and Salts of Phosphorus. There are three 
phosphoric acids, orthophosphoric (H 3 PO 4 ), metaphos- 
phoric (HPO 3 ), and pyrophosphoric (H 4 P 2 O 7 ). Phos- 
phorous acid (H 3 PO 3 ) and hypophosphorous acid (H 3 PO 2 ) 
are important compounds. 

Orthophosphoric Acid is a by-product in the manufacture of phos- 
phorus from bone ash (see above) ; it may be made by oxidizing red 
phosphorus with nitric acid, or by dissolving phosphorus pentoxide 
in hot water, thus 

PA + 3H 2 2H 3 P0 4 

Phosphorus Pentoxide Orthophosphoric Acid 

It is a white, crystalline deliquescent solid. 

Metaphosphoric Acid is formed by heating orthophosphoric acid to a 
high temperature, thus 

H 3 P0 4 HP0 3 + H 2 / 

Orthophosphoric Acid Metaphosphoric Acid 
It may be formed by dissolving the pentoxide in cold water, thus 
P 2 O 5 -f H 2 O = 2HPO 3 . 

At ordinary temperature it is a glassy solid, and is called glacial phos- 
phoric acid. It dissolves readily in water, and the solution changes 
into orthophosphoric acid slowly in the cold, rapidly when boiled. 

Phosphorus, Arsenic, Antimony, Bismuth. 269 

Pyrophosphoric Acid is formed by heating orthophosphoric acid to 
200 -300 C., thus 

2H 3 PO 4 = H 4 P 2 O 7 + H 2 O 

Orthophosphoric Acid Pyrophosphoric Acid 

A sodium salt of the ortho-acid is usually used. It may also be formed 
thuS ~ PA + 2H 2 = H 4 P 2 7 . ^ 

This acid is an amorphous, glassy (but sometimes crystalline) solid. 
It is readily soluble in water, and its solution behaves like metaphos- 
phoric acid. 

Orthophosphoric acid is tribasic, and its salts, the phosphates, are 
numerous. The most important is the normal calcium salt, Ca 3 (PO 4 ) 2 . 
Hydrogen disodium phosphate (HNa 2 PO 4 ) is the commercial sodium 
phosphate. This salt and hydrogen sodium ammonium phosphate, or 
microcosmic salt (HNa(NH 4 )PO 4 ), are used in chemical analysis. The 
" acid phosphate " sold as a beverage is a solution of one or more acid 
calcium phosphates (HCaPO 4 and H 4 Ca(PO 4 ) 2 ). Metaphosphates are 
formed by heating primary or (mono-) sodium phosphates, thus 

H 2 NaPO 4 NaPO 3 + H 2 O 

Primary Sodium 

Sodium Phosphate Metaphosphate 

Pyrophosphates are formed by heating secondary (or di-) phosphates, 
thus 2HNa,P0 4 Na 4 P 2 O 7 + H 2 O 

Disodium Phosphate Sodium Pyrophosphate 

Hypophosphites are produced by treating phosphorus with alkalies. 
They are often used as medicines. 

Other Compounds of Phosphorus. Phosphine (PH 3 ) is analogous 
to ammonia (NH 3 ), though it is not alkaline. It is made by heating 
sodium (or potassium) hydroxide with phosphorus. It is poisonous, 
has a disagreeable odor, and burns in the air, owing to the presence of 
an inflammable compound of phosphorus and hydrogen. Phosphine 
itself does not burn. It combines with other substances, forming 
phosphonium compounds, which are analogous to ammonium com- 
pounds, e.g. 

PH 3 + HI PH 4 I >/ 

Phosphine Hydriodic Acid Phosphonium Iodide 

270 Descriptive Chemistry. 

Phosphorus Trichloride (PCI.,) is a disagreeable smelling liquid, made 
by the combustion of dry chlorine and phosphorus ; and phosphorus 
pentachloride (PC1-) is a greenish solid made by passing chlorine into 
a vessel containing the trichloride. 

Matches. Phosphorus is chiefly used in the manufac- 
ture of matches. Soft wood is cut by machinery into the 
desired shape. The cards or sticks are fixed in a frame, 
and one end is first dipped into melted sulphur or paraffin 
and then into the phosphorus mixture. The latter consists 
usually of different proportions of phosphorus, manganese 
dioxide, glue, and a little coloring matter. Manganese di- 
oxide may be replaced by other oxidizing agents. These 
matches are the ordinary friction or sulphur kind. By 
rubbing them on a rough surface enough heat is gener- 
ated to cause the phosphorus to unite with the oxygen of 
the oxidizing agent, and the heat thereby produced sets 
fire to the sulphur or paraffin, and this in turn kindles the 
wood. Since these matches are poisonous, and liable to 
take fire, their manufacture has been prohibited in some 
countries (e.g. Switzerland and the Netherlands). Safety 
matches, which replace them, contain no yellow phos- 
phorus. The head of this kind is usually a colored mix- 
ture of antimony sulphide, potassium chlorate, and glue; 
while the surface upon which the match must be rubbed to 
light is coated with a mixture of red phosphorus, glue, and 
powdered glass. Matches are made by machinery, several 
million being produced in one day. 

Relation of Phosphorus to Life. Phosphorus is essen- 
tial to the growth of plants and animals. Plants take 
phosphates from the soil and store up the phosphorus 
compounds, especially in their fruits and seeds. Animals 
eat this vegetable matter, assimilate the phosphorus com- 
pounds, and deposit them in the bones, brain, and nerve 

Phosphorus, Arsenic, Antimony, Bismuth. 271 

tissue. Bones contain about 60 per cent of calcium phos- 
phate. Part of the combined phosphorus consumed by 
animals is rejected by them, and often finds its way back 
into the soil. 

The constant removal of phosphates by plants would soon exhaust 
the soil. Hence phosphorus is restored to the soil in the form of natu- 
ral or artificial fertilizers. Natural fertilizers are (i) stable refuse, 
which always contains some of the phosphates from the food originally 
fed to the animals ; (2) guano, which is the dried excrement and carcasses 
of the sea birds that once lived in vast numbers in Peru and Chili ; and 
(3) phosphate slag, which is a phosphorus by-product obtained in manu- 
facturing steel. These and bones are ground and spread upon the soil. 
Artificial fertilizers are made from phosphate rock. This occurs in large 
beds in South Carolina, Tennessee, and Florida, which yield about a 
million tons a year. It consists of the hardened remains of land and 
marine animals, and is mainly tricalcium phosphate (Ca 3 (PO 4 ) 9 ). It is 
insoluble in water, and must be changed into the soluble monocalcium 
salt (H 4 Ca(PO 4 ) 2 , so that it can be evenly distributed through the soil 
and easily taken up by plants. This soluble salt is called " superphos- 
phate of lime." When phosphate rock is treated with sulphuric acid, 
the changes involved may be written thus 

Ca 3 (PO 4 ) 2 + '2H 2 SO 4 = H 4 Ca(PO 4 ) 2 + 2CaSO 4 
Tricalcium " Superphosphate Calcium 

Phosphate of Lime " Sulphate 

Ca 3 (P0 4 ) 2 + 3H 2 S0 4 = 2H 3 P0 4 + 3 CaSO 4 

Phosphoric Acid 

Ca 3 (PO 4 ) 2 + H 2 SO 4 = H 2 Ca 2 (PO 4 ) 2 + CaSO 4 

Dicalcium Phosphate 

The aim is to convert the crude phosphate rock into "superphos- 
phate," but the three reactions usually occur. The product is ground, 
dried, and packed in bags for the market. On standing, it may undergo 
" reversion," i.e. the " superphosphate " and phosphoric acid may form 
insoluble phosphates, thus making the fertilizer less valuable. Some- 
times " superphosphate " is mixed with compounds of nitrogen and of 
potash to produce a complete fertilizer. 

272 Descriptive Chemistry. 


Occurrence. Arsenic is found free in nature, but it 
usually occurs combined with sulphur or a metal, or with 
both. The common arsenic ores are realgar (As 2 S 2 ), 
orpiment (As 2 S 3 ), arsenic pyrites or mispickel (FeSAs). 
Arsenic trioxide or arsenolite(As 2 O 3 ) is also found. Small 
quantities of arsenic occur in many ores. 

The United States annually imports over 6,000,000 pounds of arsenic 
and its compounds, mainly from England and Germany. 

Arsenic is prepared in the laboratory by heating a mixture of arse- 
nious oxide and charcoal in a glass tube. The change is represented 

2As 3 O 8 -f 6C As 4 + 6 CO 

Arsenious Oxide Carbon Arsenic Carbon Monoxide 

On a large scale it is extracted from its ores either by the above method 
or by roasting arsenic pyrites (FeSAs) in the absence of oxygen. 

Arsenic has marked properties. It is a brittle, steel-gray solid. A 
freshly broken piece has a metallic luster, which disappears slowly in a 
moist atmosphere. It tends to crystallize. The specific gravity is from 
5.62 to 5.96. Heated in the air, it volatilizes without melting, and the 
vapor has an odor like garlic. At about 180 C. it burns in the air with 
a bluish flame, forming the white oxide (As 2 O 3 ). Arsenic molecules, 
like those of phosphorus, contain four atoms. In some respects arsenic 
resembles both metals and non-metals. It is used to harden the lead 
which is made into shot. 

Arsenious Oxide or Arsenic Trioxide, As 2 O 3 , is the 
most important compound of arsenic, and is often called 
simply " arsenic" or "white arsenic." It is found free in 
nature, but is usually manufactured by roasting arsenic 
ores. There are two common varieties, a white, granular 
powder and an amorphous, glasslike solid. It has no odor, 
a faint, metallic taste, dissolves slightly in cold water, but 
readily in hot hydrochloric acid. Arsenic trioxide is a 

Phosphorus, Arsenic, Antimony, Bismuth. 273 

rank poison. The antidote is fresh ferric hydroxide, which 
is made by adding ammonium hydroxide to a ferric salt, 
e.g. ferric chloride. Small doses (2 to 3 grains) are usually 
fatal, but by habitual use the system appropriates larger 
doses without ill effects. Workmen in arsenic factories 
often accidentally swallow with impunity quantities which 
would ordinarily prove fatal. It is used for making pig- 
ments for green paints, for fly and rat poison, in mak- 
ing glass, arsenic compounds, in calico printing, and in 
preserving skins. As a medicine it is used to purify the 

Other Arsenic Compounds. The native mineral orpiment (As 2 S 3 ) 
is used in making a- yellow paint, and realgar (As 2 S 2 ) a red paint. 
Scheele's green is chiefly copper arsenite (HCuAsO 3 ), and was formerly 
used to make a cheap green paint and to color wall paper. The com- 
plex arsenic compound Paris green is a light green powder ; owing to 
its poisonous character it is used to exterminate potato bugs and other 
insects. Arsenic forms acids analogous to the acids of phosphorus, 
though they are less important. The salts sodium arsenate (HNa 2 AsO 4 ) 
and arsenite (NaAsO 2 ) are used in dyeing. The formation of the yel- 
low sulphide (As 2 S 3 ) by passing hydrogen sulphide into an arsenic 
solution containing hydrochloric acid is the usual test for arsenic. 

Marsh's Test for Arsenic. Arsenic itself is not poisonous, but its 
compounds are among the most poisonous substances known. For- 
tunately, combined arsenic is easily detected by a simple method, called 
Marsh's test. An apparatus for generating hydrogen is provided with 
a hard glass horizontal delivery tube, narrowed in places and drawn to 
a point. Pure zinc, pure dilute sulphuric acid, and the arsenic solution 
are put in the generator. Hydrogen and gaseous hydrogen arsenide 
(or arsine (AsHo) ) are formed. If this mixture is lighted at the end 
of the delivery tube, metallic arsenic is deposited as a black coating on 
cold porcelain held in the flame ; or if the tube is heated in front of a 
narrow place, arsenic is deposited at this point. This deposit dissolves 
in sodium hypochlorite solution, but a deposit of antimony, similarly 
produced, does not dissolve. By this delicate test the merest trace of 
arsenic is readily and positively detected. 

274 Descriptive Chemistry. 


Occurrence of Antimony. Small quantities of free anti- 
mony are found. The most common ore is stibnite (Sb 2 S 3 ), 
which occurs in Japan, Austria-Hungary, France, Algeria, 
Italy, Mexico, and Turkey. Large deposits in California 
and Nevada are now utilized, about 3,000,000 pounds being 
annually produced. 

Stibnite was known in the fifteenth century. The Latin name of 
antimony is stibium, from stibnite, which gives the symbol of the 
element, Sb. 

Antimony is prepared on a large scale by two methods. In one the 
sulphide is roasted, and the oxide thus formed is reduced with charcoal. 
Equations representing the main changes are 

2Sb 2 S 3 + 9O 2 = 2SbO 3 4 6 SO 2 
Antimony Sulphide Oxygen Antimony Oxide Sulphur Dioxide 

2Sb 2 O 3 + 3C 4Sb + 3CO 2 

The other method consists in heating the sulphide with iron, the equation 
for the chemical change being 

Sb 2 S 3 + 3Fe = 2Sb -f 3 FeS 

Antimony Sulphide Iron Antimony Iron Sulphide 

Antimony has interesting properties. It is a silver white, crystal- 
line, brittle solid. Its specific gravity is 6.7. At ordinary temperatures 
antimony does not tarnish in the air, but when heated, it burns with a 
bluish flame, forming the white, powdery antimony trioxide (Sb 2 O.,). 
Powdered antimony burns brilliantly when added to chlorine, bromine, 
or iodine. Nitric acid oxidizes it, and aqua regia dissolves it. Anti- 
mony melts at about 450 C. It expands on cooling, and is therefore 
one constituent of type metal (see Alloys of Lead). 

Compounds of Antimony. Antimony forms stibine (SbH 3 ), which 
is analogous to ammonia (NH 3 ) and arsine (AsH 3 ), pyro- and meta- 
acids, the oxides, Sb 2 O 3 and Sb 2 O-, and halogen compounds. It also 
forms complex compounds in which antimony acts as a metal. Tartar 
emetic is potassium antimonyl tartrate (KSbO .C 4 H 4 O 6 ). It is used as 
a medicine and as a mordant in dyeing cotton. Antimony trisulphide 

Phosphorus, Arsenic, Antimony, Bismuth. 275 

(Sb.,S 3 ) is a reddish solid, formed by passing hydrogen sulphide gas 
into a solution of antimony the test for antimony. The sulphide is 
used in making the red rubber tubing and stoppers used in the labora- 
tory. The chlorides (SbCl 3 and SbCl 3 ) are formed by the action of 
chlorine upon the metal ; with water they form the white solids called 
oxychlorides, e.g. SbOCl. The formation of antimony oxychloride is 
sometimes used as a test for antimony, but the more common test is 
the formation of the reddish orange sulphide (Sb 2 S 3 ). 


Bismuth is usually found in the native state, though it is 
not abundant nor widely distributed. The oxide (Bi 2 O 3 ), or 
bismite, the carbonate ((BiO) 2 CO 3 .H 2 O), or bismutite, and 
the sulphide (Bi 2 S 3 ), or bismuthinite, are the common ores. 
The world's supply comes from Saxony. 

Bismuth is prepared from the native metal by melting it on an 
inclined plate and allowing it to drain away from the solid impurities. 
Sometimes the sulphide is roasted, and the resulting oxide is reduced 
with charcoal, as in the case of antimony. 

Bismuth has characteristic properties. It is a grayish white metal 
with a reddish tinge. Like antimony, it is very brittle. It does not 
tarnish in dry air, but it grows dull in moist air ; and when heated in 
air it burns with a bluish flame, forming the yellowish oxide (Bi 2 O 3 ). 
Its specific gravity is about 9.9. Hydrochloric acid does not readily 
attack it, but nitric acid converts it into a nitrate, and hot sulphuric acid 
into a sulphate. 

Bismuth melts at about 270 C. But a mixture of bismuth, lead, and 
tin melts at a low temperature. For example, Newton's metal melts at 
95 C. and Rose's metal at 100 C. ; while Wood's metal, which con- 
tains cadmium, melts at only 66 C.-yi C. These metallic mixtures 
are called fusible metals. They are used in making casts of wood 
cuts; but more often (i) as safety plugs in steam boilers to prevent 
explosions, (2) as a fuse in electrical apparatus to prevent a short cir- 
cuit, and (3) to hold in place fireproof doors and the valves in the 
automatic sprinkling apparatus now placed in large buildings. 

Compounds of Bismuth. Bismuth forms no compound with hydro- 
gen. There are three oxides. Bismuth trioxide (Bi 2 O 3 ) is yellowish, 

2j6 Descriptive Chemistry. 

the pentoxide (Bi 2 O 5 ) is orange red, and the dioxide (Bi 2 (X) is black. 
Bismuth trioxide is used to fix the gilding on porcelain. The trichloride 
(BiCl t3 ) is formed by the action of chlorine upon bismuth, or by treat- 
ing bismuth with aqua regia. With an excess of water the trichloride 
forms the oxychloride (BiOCl), which is a pearl-white powder, insoluble 
in water. The formation of the oxychloride is the usual test for bis- 
muth. Bismuth, being a metal, forms hydroxides (Bi(OH) 3 and 
BiO.OH). Normal bismuth nitrate (Bi(NO. ? ) ;i ), treated with hot 
water, forms basic bismuth nitrate (Bi(OH) 2 NCX or BiONO 3 ). The 
latter, often called subnitrate of bismuth, is a white powder used as a 
medicine for dyspepsia and as a cosmetic. 


1 . What is the symbol and atomic weight of phosphorus ? Give a 
brief history of this element. Why is it so named ? 

2. Discuss the occurrence of phosphorus. 

3. Describe the manufacture of phosphorus (a) from a .phosphate 
and sulphuric acid, and (b) by the electric method. How is it 
purified ? 

4. Summarize the properties of (a) ordinary phosphorus, and 
() red phosphorus. 

5. Describe briefly (a) the oxides of phosphorus, (b} orthophos- 
phoric acid, (c) metaphosphoric acid, (d} pyrophosphoric acid, (e) phos- 
phine, (/) the phosphorus chlorides. 

6. What is (a) tricalcium phosphate, (b} microcosmic salt, (c} " acid 
phosphate " ? 

7. Describe the manufacture of (a) sulphur matches, and (b} safety 

8. Discuss the relation of phosphorus to life. 

9. What is a fertilizer ? Name three natural fertilizers. Describe 
the manufacture of artificial fertilizer. What is a complete fertilizer ? 

10. What is the symbol and atomic weight of arsenic ? 

1 1 . Name several ores of arsenic. With what metals is arsenic often 
associated ? 

12. Describe the preparation and state the properties of the arsenic. 

13. What is the formula of arsenic trioxide ? By what other names 
is it known ? Summarize its properties. For what is it used ? What 
is the antidote for arsenic poisoning ? 

Phosphorus, Arsenic, Antimony, Bismuth. 277 

14. What is (a) Paris green, (6) orpiment, (V) realgar ? For what 
is each used ? 

15. Describe Marsh's test for arsenic. 

1 6. What is the symbol and atomic weight of antimony ? 

17. In what forms does antimony occur and where is it found ? De- 
scribe its preparation. State its chief properties. 

18. What is tartar emetic ? For what is it used ? 

19. Describe the test for antimony. 

20. What is the symbol of bismuth ? How does it occur and where 
is it found ? Describe its preparation. State its properties. 

21. State the relation of bismuth hydroxide to bismuth subnitrate. 
Describe the latter. 


1. Calculate the percentage composition of (a) sodium phosphate 
(Na 3 PO 4 ), () dihydrogen phosphate (H 2 NaPO 4 ), (V) disodium phos- 
phate (HNa 2 PO 4 ), O/) microcosmic salt (HNaNH 4 PO 4 ). 

2. How much phosphorus is needed to remove the oxygen from a 
liter of air ? (Assume (i)2P + 5O = P 2 O 5 and (2) air is 20 per cent 

3. How much phosphorus is there in a ton (2000 Ib.) of bone ash 
(Ca 3 (P0 4 ) 2 )? 

4. If a skeleton weighs 25 Ib. and contains 60 per cent calcium 
phosphate, how much phosphorus does it contain ? 

5. What is the weight of a cylindrical stick of ordinary phosphorus 
10 cm. long and 15 mm. in diameter ? (Suggestion. What is the spe- 
cific gravity of phosphorus ?) 

6. Calculate the percentage composition of (a) orpiment (As 3 S 3 ), 
($) realgar (As S 2 ), (^) white arsenic (As 2 O 3 ). 

7. What is the weight of a piece of antimony 25 cm. long, 15 cm. 
wide, and 2 mm. thick ? 


Introduction. The elements studied thus far are chiefly 
non-metals. Metals, however, have been mentioned, and 
many of their properties have been discussed. It is the 
purpose of the present chapter to review these properties 
and prepare the way for a fuller treatment of the metals. 

Metals and Non-metals. Many years ago the chem- 
ical elements were divided into two classes, called metals 
and non-metals. The division was based largely on the 
physical properties of the elements. The opaque, lustrous, 
more or less heavy, hard, ductile, malleable, tenacious 
solids were called metals. All gases and the solids such 
as carbon, sulphur, phosphorus, and iodine were called 
non-metals. No such sharp dividing line, however, can 
be drawn between metals and non-metals. Some, of 
course, have pronounced properties, like the non-metal 
sulphur and the metal iron. These are typical. But a 
few have variable properties. Sometimes they act as 
metals and at other times as non-metals. Antimony and 
arsenic belong to this border-line class ; they are sometimes 
called the metalloids. The classification into metals and 
non-metals is no longer accurate, but it is very convenient. 
The use in common life of the words metallic and metal 
seldom leads to confusion. 

Properties of Metals. The physical properties of 
metals are familiar, though variable between wide limits. 


Metals. 279 

All have a metallic luster, i.e. the marked property of 
reflecting light from their polished or untarnished surfaces. 

All are opaque except very thin films of gold. The 
color of many is white, though the tint varies. Thus 
silver, sodium, aluminium, mercury, magnesium, iron, and 
tin are nearly pure white, and bismuth is reddish white. 
Copper is the only red metal, and gold the only yellow 
one, which is an element. Most metals are malleable and 
ductile, i.e. they may be hammered or rolled into sheets 
and drawn into wire. Gold, copper, silver, iron, platinum, 
and aluminium possess both these properties to a marked 
degree ; while lead, zinc, and tin are very malleable though 
not so ductile. Antimony and bismuth are brittle. The 
hardness of metals varies. At the ordinary temperature 
mercury is a liquid, sodium and lead can be cut easily with 
a knife, and so on through the list up to iridium, which is 
as hard as steel. In specific gravity, which was once 
thought must very high, the metals range between lithium, 
which has the specific gravity 0.585, and osmium, which has 
the specific gravity 22.48. Sodium and potassium also are 
lighter than water, while magnesium has the specific grav- 
ity 1.75, and aluminium 2.58. Metals are good conductors 
of heat and electricity. They also vary in this property. 
Silver, copper, and aluminium are the best conductors, and 
have therefore many practical applications. Bismuth is 
the poorest conductor. 

The distinctive property of metals is not physical, but 
chemical. Metals form oxides which combine with water 
to produce bases. Metals are the characteristic elements 
of bases. On the other hand, non-metals form acid-pro- 
ducing compounds. 

Occurrence of Metals. Only a few metals are found 
free in the earth's crust, and these are seldom pure. Of 

280 Descriptive Chemistry. 

the six metals known to the ancients, gold, copper, silver, 
tin, iron, and lead, only gold and copper are found free. 
The solid elements and their compounds which occur in 
the earth's crust are called minerals. And those minerals 
from which metals can be profitably extracted are called 
ores. The most abundant classes of ores are oxides, sul- 
phides, carbonates, and hydroxides. Lead, zinc, mercury, 
and silver sulphides are abundant. Besides native copper, 
the sulphide and carbonate are found. Iron occurs as 
oxide, carbonate, hydroxide, and sulphide. Many ores 
contain arsenic. Some ores are very complex. 

Preparation of Metals. The series of operations by 
which useful metals are extracted from their ores is called 
metallurgy. It includes preliminary treatment, smelting, 
electrolysis, refining, and other operations necessary to 
change the ore into a metal ready for manufacture into 
useful articles. The object of the preliminary treat- 
ment is to prepare the ore for smelting or for a similar 
operation by which the metal is obtained in a state 
adapted for further purification or refining. The ore as it 
comes from the mine is usually mixed with earthy matter 
or rock called gangue. This impurity is removed by me- 
chanical or chemical processes, sometimes by both. The 
mechanical process illustrates one kind of preliminary treat- 
ment. The ore is first crushed in a stamp mill. This is a 
huge, heavy mortar and pestle. The pestle or stamp falls 
repeatedly upon the ore, which is slowly fed into the mortar 
or die. A current of water (or air) forces the fine particles 
out of the mortar through a sieve. The lighter particles of 
the impurities are washed away, and the metallic grains 
are extracted by another mechanical operation, though 
chemical processes are frequently employed, especially 
with inferior ores. This separation of the valuable part 

Metals. 281 

of the ore from the gangue, and reducing it to a smaller 
bulk is often called ore dressing or concentration. Copper 
is extracted from the Lake Superior ores mainly by this 
method of preliminary treatment. 

Gold and silver ores are treated this way, and then ex- 
tracted from the slime by mercury. The latter operation 
is called amalgamation. The most common method of 
extracting metals from their ores is by smelting. The 
process varies with the kind and composition of the ore. 
Essentially, it consists in heating a mixture of the ore and 
coke (or coal) in a furnace. The ores used must, as a rule, 
be oxides. Sulphides, hydroxides, and carbonates are first 
roasted or calcined to convert them into oxides. The 
essential chemical change in smelting is a reduction of the 
oxide by the carbon. The carbon and oxygen unite and 
pass off as a gas, leaving the metal to run out at the bot- 
tom. Limestone, or a similar substance, called a flux, is 
added to the mixture, if necessary, to facilitate the melting 
and to assist in removing the impurities as a glassy sub- 
stance, called slag. The operation is conducted in differ- 
ent kinds of furnaces. Iron, for example, is smelted in a 
huge upright furnace called a blast furnace (Fig. 72), 
because a current of air is forced through the melted mass 
to facilitate the fusion and chemical changes. In such a 
furnace the fuel and ore are in direct contact. When this 
is undesirable, the reverberatory furnace is used (Fig. 54). 
As the figure shows, in this furnace the flame is reflected 
or reverberated upon the ore under treatment. In this 
kind of furnace the ore may be oxidized or reduced with- 
out coming in contact with the fuel. Some ores demand 
special methods, which will be described in connection 
with these metals. Electrolysis is used to extract some 
metals, especially aluminium. Other metals, notably 


Descriptive Chemistry. 

FIG. 54. Reverberatory furnace. Tue tire 

copper, are purified by 
electrolysis. A few met- 
als are extracted by a 
wet process. That is, 
the ores are dissolved, 
and the metal is precipi- 
tated by adding some 
substance or by elec- 

i ivj. 5*j- - AV<- v^i u\,i aiwi j mi lit L\^... i nt; me -I . T*!. * X" " 

burns on the grate, G, and the long flame trolySlS. 1 huS, interior 

which passes over the bridge, E, is reflected Pjold Ores are dissolved 

down by the sloping roof upon the contents . , 

of the furnace. Gases escape through /. The by treatment With potas- 

charge, which rests upon B, does not come sium CV^nide and the 

in contact with the fuel, but is oxidized or . . 

reduced by the flame. gold IS then precipitated 

by zinc. 

Alloys are mixtures or compounds of two or more 
metals. Some fused metals mix in all proportions, while 
others seem to form definite compounds. The properties 
of alloys vary with the constituents and their properties. 
Some alloys, especially those of copper and of lead, have 
many industrial uses. Alloys in which mercury is a con- 
stituent are called amalgams. 


1. Define the terms metal and non-metal as they are ordinarily used. 
Name six or more examples of each class. Define and illustrate the 
term metalloid. Why is this classification inaccurate? 

2. State the familiar physical properties of metals. Define (a) 
metallic luster, (b} malleable, (c) ductile, (d) specific gravity. 

3. How does the color of metals differ from their luster? Name 
five metals which are white. What color has (a) gold, (6) copper, (c) 
zinc, (a) lead, (e) iron? 

4. What metals are brittle? Malleable? Soft? Hard? Heavy? 
Light? What metals conduct electricity well? 

5. What is the distinctive chemical property of metals? Of nor>' 
metals? Illustrate your answer. 

Metals. 283 

6. What metals are often found free in nature? Define and illus- 
trate the terms (a) mineral, and (b) ore. What are the most abundant 
classes of ores? 

7. What metals occur abundantly as (a) sulphides, (b) oxides, (V) 
carbonates ? 

8. Define metallurgy. W T hat general operations does it include? 

9. What is the object of the preliminary treatment of ores? How 
is it accomplished mechanically? Define (a) gangue, (<) concentra- 
tion, (c) amalgamation. What metal is often extracted (a) mechanic- 
ally, (b) by amalgamation ? 

10. Define smelting. What fundamental chemical change does it 
usually involve? Define and illustrate (a) calcination, ($) flux, (c) 

n. Describe (a} a reverberatory furnace, and (b) a blast furnace. 
What is their essential difference ? 

12. What is the wet process of extracting ores? 

13. What are (a) alloys, (b) amalgams? 


1. What is the specific gravity of gold, if a piece weighs 4.676 gm. 
in air,' and loses 0.244 gn. when weighed in water? (Note. Specific 
gravity equals the weight in air divided by the loss of weight in water.) 

2. A piece of aluminium weighs 150 gm. in air and 75 gm. in water. 
What is its specific gravity? 

3. A piece of iron weighs 292.8 gm. in air and 255.3 gm. in water. 
What is its specific gravity? 

4. A piece of copper weighing 50 gm. in air lost 5.6 gm. when 
weighed in water. What is its specific gravity? 

5. A piece of lead pipe weighs 158.9 gm. in air and 144.9 & m - ^ n 
water. Calculate the specific gravity. 


Introduction. Sodium and potassium, and the rare 
elements lithium, rubidium, and caesium, form a natural 
group, known as the alkali metals. The different elements 
and their corresponding compounds resemble each other 

Sodium and potassium were discovered by Sir Humphry Davy in 
1807 by the electrolysis of their hydroxides. Bunsen, by means of the 
spectroscope, discovered lithium in 1855, caesium in 1860, and rubidium 
in 1861. 


Occurrence. Sodium is not found free. Sodium chlo- 
ride and sodium nitrate are the most abundant compounds. 
Many rocks, plants, and mineral waters contain combined 
sodium. About 2.5 per cent of the earth's crust is sodium. 

The symbol of sodium, Na, is from the Latin word natrium, which 
in turn comes from the Greek word natron, an old name of sodium 

Preparation. Sodium is now manufactured on a large 
scale by the electrolysis of fused sodium hydroxide. This 
method was used by Davy in 1807 to isolate sodium, but 
its commercial success was only recently made possible by 
Castner. Figure 55 is a sketch of the apparatus used. 
The body of the steel cylinder, S, rests within a heated 
flue. Hence the sodium hydroxide is solid in the neck, B, 
and serves to protect the joint made by the iron cathode, 





Sodium, Potassium, and Lithium. 


C, and the crucible. A, A is the 
iron anode. A collecting pot, 
P, dips into the molten caustic 
soda. As the electrolysis pro- 
ceeds, the sodium formed at 
C collects in P, and a wire 
gauze, G, G, keeps it from mix- 
ing with the caustic soda. 
The sodium is ladled out at 
intervals from P. The hy- 
drogen, which is liberated, 
accumulates also in P and 
prevents the sodium from oxi- 
dizing. The hydrogen some- 
times escapes and explodes. FIG. 55. Apparatus for the manu- 
facture of sodium by the electrolysis 

Sodium was formerly manufactured O f sodium hydroxide, 
by two methods, (i) Sodium car- 
bonate and carbon heated to a high temperature change thus 

Na 2 C0 3 
Sodium Carbonate 

2C = 2Na 
Carbon Sodium 

+ 3 CO 

Carbon Monoxide 

The mixture was heated in iron retorts, and the sodium vapor, in pass- 
ing through a flat iron receiver, condensed to a liquid, which was col- 
lected under paraffin or mineral oil. (2) The other chemical method, 
devised by Castner in 1886, consisted essentially in heating sodium 
hydroxide with a mixture of iron and carbon. Probably iron carbide 
was the essential reducing agent, and the change might be represented 

6NaOH + FeC 2 = 2 Na + Fe + 2 Na,CO, + 3 H 2 
Sodium Hy- Iron Car- Sodium Iron Sodium Car- Hydrogen 




Properties. Sodium is a silver-white metal. It is so 
soft that it may be easily molded with the fingers and 
cut with a knife. It floats upon water, since its specific 

286 Descriptive Chemistry. 

gravity is only 0.98. Heated in the air, it melts at 96 C, 
and at a higher temperature it burns with a brilliant yellow 
flame, forming the oxides Na 2 O and Na 2 O 2 . This intense 
yellow color is characteristic of sodium and is the usual 
test for the element (free or combined). In moist air the 
bright surface quickly tarnishes, and sodium as usually 
seen has a brownish coating. It is, therefore, kept under 
kerosene or a liquid free from water. It decomposes 
water at ordinary temperatures, liberating hydrogen and 
forming sodium hydroxide, thus 

Na + H 2 O - NaOH + H 
Sodium Water Sodium Hydroxide Hydrogen 
If held in one place upon water by filter paper, enough 
heat is generated to set fire to the hydrogen, which burns 
with a yellow flame, owing to the presence of volatilized 
sodium (see Interaction of Sodium and Water, Chapter V). 
If melted sodium is put into chlorine, the two elements 
combine with a brilliant flame, forming sodium chloride. 
It was in this way that Davy, in 1810, proved that com- 
mon salt is really nothing but sodium chloride. It combines 
directly with the other halogens. 

A molecule of sodium contains only one atom. 

Sodium is used in the laboratory to extract water from alcohol and 
ether and to prepare organic compounds. Large quantities are con- 
sumed in the manufacture of sodium peroxide (Na a O 2 ) and sodium 
cyanide (NaCN). Its power to reduce oxides gives it limited use in 
preparing certain metals, e.g. magnesium. 

Sodium Chloride, NaCl, is the most important compound 
of sodium. It is familiar under the name of salt or com- 
mon salt. The presence of salt in the ocean, in lakes and 
springs, and in the soil is mentioned in the oldest histori- 
cal records. It is one of the most abundant substances. 
The sources of salt are sea water, rock salt, and brines. 

Sodium, Potassium, and Lithium. 287 

Preparation of Salt. Sea water contains nearly 4 per cent of salts, 
and three fourths of this amount is sodium chloride, (i) In warm 
countries, as on the shores of the Mediterranean Sea, shallow ponds of 
sea water near the shore are evaporated by exposure to the sun and 
wind, and the salt is collected. (2) In some regions sea water is first 
concentrated by allowing it to trickle over heaps of brush and then 
evaporated to crystallization in shallow pans. (3) In cold countries, 
as on the shores of the White Sea in Russia, sea water is allowed to 
freeze and the ice is removed. The ice contains no salt, so the opera- 
tion is repeated until the remaining liquid becomes strong enough to 
evaporate profitably over a fire. (4) Deposits of salt are found in 
many parts of the globe, the most important being in England, Austria- 
Hungary, and Germany. In these regions and some parts of the 
United States, the salt is mined and purified like other minerals. This 
variety is coarse and often impure, and is largely used in curing meat 
and preserving hides. (5) Most of the salt produced in the United 
States is obtained from natural or artificial brines, i.e. from strong solu- 
tions of salt. Artificial brines are made by forcing water into salt de- 
posits. Brines are obtained in New York, Michigan, Kansas, Ohio, 
West Virginia, California, Utah, and Louisiana. They are evaporated 
in vats by the sun's heat or by heating in kettles or pans. 

All these methods give a product containing as impurities salts of 
sodium, calcium, and magnesium, which are largely removed by further 
special treatment. The dampness of salt is due mainly to the magne- 
sium chloride it contains (see Deliquescence, Chapter IV). * 

Properties and Uses of Salt. Salt is soluble in water, 
100 gm. of water dissolving about 36 gm. of salt at o C., 
and 40 gm. at 100 C. It crystallizes in cubes. This sub- 
stance is an essential ingredient of the food of man and 
animals. Besides its universal domestic use, enormous 
quantities are consumed in the preparation of many so- 
dium compounds, particularly sodium carbonate (see below), 
of hydrochloric acid and bleaching powder. In 1902 the 
United States produced nearly 3,000,000 tons of salt, and 
imported over 200,000 tons. This is about the average 

288 Descriptive Chemistry. 

Sodium Carbonate, Na 2 CO 3 , is next to sodium chloride 
in importance. Small quantities of hydrated sodium car- 
bonates are found in Egypt, Russia, and in California and 
Nevada. Formerly it was obtained from the ashes of 
marine plants, but sodium chloride is now the source. 
The manufacture of sodium carbonate is one of the most 
extensive chemical industries. Two processes are used, 
the Leblanc and the Solvay. 

The Leblanc Process has three steps, (i) Sodium chloride is 
changed into sodium sulphate by sulphuric acid, the two equations for 
the changes being 

2NaCl + H 2 SO 4 = HNaSO 4 + HC1 + NaCl 
Sodium Sulphuric Acid Sodium Hydrochloric Sodium 

Chloride Acid Sulphate Acid Chloride 

HNaSO 4 + NaCl Na,SO 4 + HC1 

Sodium Sulphate 

This operation is called the "salt cake process 11 ; the impure prod- 
uct, called " salt cake," contains about 95 per cent of sodium sulphate. 
The hydrochloric acid is a by-product (see Hydrochloric Acid) . (2) 
The sodium sulphate is changed into sodium carbonate by heating the 
"salt cake" with coal and limestone, the main changes being repre- 
sented by the equations 

Na 2 S0 4 + 2 C Na 2 S + 2 CO, 

Sodium Sulphate Carbon Sodium Sulphide Carbon Dioxide 

Na 2 S + CaCO 3 Na 2 CO 3 + CaS 

Sodium Lime- Sodium Calcium 

Sulphide stone Carbonate Sulphide 

This operation is called the "black ash -process." The product is a 
dark brown or gray porous mass, and contains, besides 37 to 45 per cent 
of sodium carbonate, considerable calcium sulphide and other impuri- 
ties. The calcium sulphide is a source of sulphur (see Sulphur). (3) 
The sodium carbonate is rapidly separated from the insoluble portions 
of the " black ash " by agitation with a small amount of cool water. 
The solution of sodium carbonate thus obtained is evaporated to crys- 

Sodium, Potassium, and Lithium. 289 

tallization, and the crude crystals are ignited. This product is known 
as soda ash, and from its solution in waiter are obtained soda crystals 
or sal soda (Na 2 CO 3 . 10 H,O). 

The Solvay Process, often called the ammonia -soda process, con- 
sists in saturating a cold concentrated solution of sodium chloride first 
with ammonia gas and then with carbon dioxide gas. The equation 
for the chemical change is . it 

NaCl + NH 3 + CO 2 = HNaCO., + NH 4 C1 
Sodium Ammonia Carbon Acid Sodium Ammonium 
Chloride Dioxide Carbonate Chloride 

The acid sodium carbonate is nearly insoluble in the cold ammonium 
chloride solution, and therefore separates. It is changed, by heating, 
into sodium carbonate, thus 

2 HNaCO, Na 2 CO 3 + CO 2 + . H 2 O 

Acid Sodium Sodium Carbon Water 

Carbonate Carbonate Dioxide 

The liberated carbon dioxide is used again, and from the ammonium 
chloride the ammonia is recovered and also used. 

Properties and Uses of Sodium Carbonate. Crystal- 
lized sodium carbonate (Na 2 CO 3 . 10 H 2 O) is often called 
alkali or soda. It loses water in the air, becoming dull 
at first and finally falling to a powder. When heated, it 
melts in its water of crystallization, and continued heating 
changes it into the white anhydrous salt (Na 2 CO 3 ). It is 
readily soluble in water, and the solution, which is strongly 
alkaline, is widely used as a cleansing agent, hence the 
name washing soda. 

Enormous quantities of sodium carbonate are used in 
the glass and soap industries, and in preparing sodium 

Sodium Bicarbonate, HNaCO 3 , is a by-product of the 
Solvay process, and it may also be prepared by treating 
crystallized sodium carbonate with carbon dioxide gas. It 
is a white powder, less soluble in water than the normal 

290 Descriptive Chemistry. 

carbonate. When heated or when mixed with an acid or 
an acid salt, sodium bicarbonate gives up carbon dioxide. 
This property early led to its use in cooking, and gives the 
names cooking soda, baking soda, or simply soda. 

Sodium bicarbonate is one ingredient of baking powder and of the 
various mixtures (except yeast) used to raise bread, cake, and other 
food. Since cream of tartar is slightly acid,- it is usually used to liber- 
ate the gas. Sour milk, which contains lactic acid, is. sometimes used 
in place of cream of tartar. When pastry is raised with soda and cream 
of tartar, the escaping carbon dioxide puffs up the dough. Hence bak- 
ing soda is often called saleratus the salt which aerates (from the 
Latin words sal, salt, and aer, air or gas). Effervescing powders, such 
as Seidlitz (or Rochelle) and soda powders, contain sodium bicarbon- 
ate in one paper and tartaric acid or one of its acid salts in the other. 
When these are mixed in water, carbon dioxide is liberated. Sodium 
bicarbonate is used as a medicine to neutralize an acid stomach. For 
example, the " soda mints " sometimes taken for this purpose are mainly 
sodium bicarbonate. 

Sodium Hydroxide or Caustic Soda, NaOH, is a white 
corrosive solid. It absorbs water and carbon dioxide 
rapidly from the air. It dissolves readily in water, with 
rise of temperature, and the solution is strongly alkaline. 
It melts easily, and is often cast into sticks for use in the 
laboratory. Immense quantities are used in making hard 
soap, paper, and dyestuffs ; in bleaching, and in refining 
kerosene oil. 

Sodium hydroxide is usually manufactured by treating crude sodium 
carbonate with calcium hydroxide. Lime is added to a boiling, dilute 
solution of soda ash, and the main change is represented thus 

Ca(OH) 2 + Na 2 CO 3 = 2 NaOH + CaCO 3 

Calcium Sodium Sodium Calcium 

Hydroxide Carbonate Hydroxide Carbonate 

The solution of sodium hydroxide is separated from the insoluble cal- 
cium carbonate, and concentrated by heating in iron kettles to the de- 

Sodium, Potassium, and Lithium. 291 

sired strength or until the mass becomes stiff. Air is then blown in or 
sodium nitrate added to oxidize sulphides to sulphates. After standing 
several hours to allow other impurities to settle, the caustic soda is put 
into iron barrels called drums. It solidifies on cooling, and the drums 
are at once sealed to keep out the air. 

Sodium hydroxide is also manufactured on a large scale 
at Niagara Falls, New York, by the electrolysis of sodium 
chloride, according to the equation 

NaCl + 


H 2 O = 

NaOH + 




+ H 


The apparatus is shown in Figure 56. The carbon anodes 
(A, A) pass into the outer compartments which contain 
brine, and the iron cathodes into the middle compartment 
which contains sodium hydroxide solution. When the cur- 

FlG. 56. Apparatus for the manufacture of sodium hydroxide by the electrolysis 
of sodium chloride. 

rent passes, chlorine is evolved at the anodes and flows out 
through pipes (not shown), and sodium is produced on 
the surface of the mercury (M ) which covers the floor 
of the whole apparatus. The sodium forms an amalgam 

292 Descriptive Chemistry. 

with the mercury, and by rocking the apparatus on the 
device, B, B, the sodium amalgam flows into the compart- 
ment, D, where the sodium is liberated by the action of the 
electric current, which passes between the cathode and 
the amalgam. The sodium reacts with the water forming 
hydrogen, which passes off through pipes (not shown) and 
sodium hydroxide, which flows into a special tank. Both 
the chlorine and sodium hydroxide are nearly pure. The 
solution of caustic soda is finally treated, if necessary, as in 
the older process. 

Sodium Sulphate, Na 2 SO 4 , is one of the products 
obtained in the manufacture of sodium carbonate (see 

In another method, sulphur dioxide, steam, and air are passed into 
hot sodium chloride. And at Stassfurt, magnesium sulphate and sodium 
chloride are allowed to interact in the cold, thus 

MgSO 4 + 2NaCl = Na.SO 4 + MgCl 2 

Magnesium Sodium Sodium Magnesium 

Sulphate Chloride Sulphate Chloride 

Sodium sulphate is a white anhydrous solid. It dissolves 
readily in water, and when a strong solution made at 30 C. 
is cooled, large transparent bitter crystals separate. They 
have the formula Na 2 SO 4 . ioH 2 O and are called Glau- 
ber's salt, from the discoverer. They lose water when 
exposed to air, and the salt continues to effloresce until it 
becomes an anhydrous powder. The crude salt is used in 
the glass and dyeing industries, and the purified salt as a 

Sodium Nitrate, NaNO 3 , is found abundantly in Chili, 
and is often called Chili saltpeter. It is a white solid, 
which becomes moist in the air. Large quantities are used 
as a fertilizer, either alone or mixed with compounds of 

Sodium, Potassium, and Lithium. 293 

potassium and of phosphorus, and for making nitric acid 
and potassium nitrate. 

The natural deposits are in a dry region near the coast and cover 
over 200,000 acres. Chili controls the industry, and exports annually 
over a million tons. The crude salt, which looks like rock salt, is puri- 
fied by crystallization into a product containing 94-98 per cent of the 
nitrate. The final mother liquor is a source of iodine (see Iodine). 

Sodium Dioxide or Peroxide, Na.,O a , is a yellowish solid. It is used 
to bleach straw and delicate fabrics. With water it liberates oxygen, 
according to the equation 

Na 2 O 2 + H 2 O + 2NaOH 

Sodium Dioxide Oxygen Sodium Hydroxide 

Miscellaneous. Sodium cyanide (NaCN) is used to extract gold 
from poor ores. Sodium monoxide (Na 2 O) is a grayish solid. The 
sodium phosphates, sodium thiosulphate, acid sodium sulphite, sodium 
silicate, and sodium tetraborate or borax have been described. 


Occurrence. This metal is not found free, but its com- 
pounds are abundant. The minerals mica and feldspar 
are silicates containing potassium. By the decay of these 
and other minerals, potassium compounds find their way 
into the soil, thence into plants and animals. Potassium 
salts are found in wood ashes, in suint, the oily substance 
washed from sheep's wool, in beet-sugar residues, and in 
the deposits in wine casks. Sea water and mineral waters 
contain potassium salts, particularly potassium chloride 
and potassium sulphate. Many potassium salts are found 
at Stassfurt. About 2.5 per cent of the earth's crust is 

The Stassfurt deposits of the salts of potassium and other metals 
are near Magdeburg, Germany. About 16 different salts make up 
the beds, which are nearly 3000 feet thick. The deposits were doubt- 
less formed by the evaporation of sea water, though the different simpler 

294 Descriptive Chemistry. 

salts interacted, forming complex ones. The most important salts 

Kainite .... KC1, MgSO 4 . 3 H 2 O. 
Carnallite . . . KC1, MgCl,, . 6 H 2 O. 
Polyhalite . . . K 2 SO 4 , Mg~SO 4 , 2 CaSO 4 . 2 H 2 O. 
Sylvite .... KC1. . 

Picromerite . . K 2 SO 4 , MgSO 4 . 6H 2 O. 

The name potassium comes from the word potash. The symbol, K, 
is from kalium, the Latin equivalent of kali, which is derived from an 
Arabic term for an alkaline substance. 

Preparation. Potassium is now obtained by the electrolysis of 
potassium hydroxide. Formerly it was manufactured, like sodium, 
by heating to a high temperature a mixture of potassium carbonate 
and carbon or of potassium hydroxide and iron carbide (see under 

Properties. Like sodium, potassium is a soft, silver- 
white metal, light enough to float upon water the specific 
gravity being 0.86. Its brilliant luster soon disappears in 
air, owing to rapid oxidation. Potassium as ordinarily 
seen is, therefore, covered with a grayish coating, and, like 
sodium, must be kept under mineral oil. It melts at 62.5 
C, and at a higher temperature burns with a violet-colored 
flame. This color is characteristic of burning potassium, 
arid is a test for the metal and its compounds. Like 
sodium, it decomposes water at ordinary temperatures, 
though more energetically. The heat evolved immediately 
ignites the hydrogen, and the melted potassium surrounded 
by a violet flame dashes to and fro upon the cold water. 
The main reaction corresponds to the equation 

K -f H 2 = KOH + H 

Potassium Water Potassium Hydroxide Hydrogen 

Potassium combines with the halogens and other ele- 
ments more vigorously than sodium, and forms analogous 

Sodium, Potassium, and Lithium. 295 

Potassium Chloride, KC1, is found native in the Stass- 
furt deposits. It is also obtained in large quantities by 
decomposing carnallite and crystallizing the potassium 
chloride from the more soluble magnesium chloride. It is 
a white solid which crystallizes in cubes and otherwise 
resembles sodium chloride. It is used chiefly to prepare 
other potassium salts, especially the nitrate and chlorate. 

Potassium bromide and potassium iodide have been described (see 
Chapter XVI). 

Potassium Nitrate, KNO 3 , is also called niter and salt- 
peter. It is formed in the soil of many warm countries 
by the decomposition of nitrogenous organic matter (see 

It is now made by mixing hot, concentrated solutions of native so- 
dium nitrate and potassium chloride, which interact thus 

NaNO 3 + KC1 KN0 3 + NaCl 

Sodium Potassium Potassium Sodium 

Nitrate Chloride Nitrate Chloride 

The sodium chloride, being less soluble, separates, and is removed. By 
evaporation, small crystals of potassium nitrate, called " niter meal, 11 are 
obtained, and further purified by recrystallization. 

Potassium nitrate is a white solid. It dissolves easily in 
cold water with a fall of temperature, and very freely in hot 
water, but it is not hygroscopic. It is crystalline, but con- 
tains no water of crystallization. The taste is salty and 
cooling. It melts at 339 C., and further heating changes 
it into potassium nitrite (KNO 2 ) and oxygen. At a high 
temperature, potassium nitrate gives up oxygen readily, 
especially to charcoal, sulphur, and organic matter. This 
oxidizing power leads to its extensive use in making gun- 
powder, fireworks, matches, explosives, and in many chemi- 
cal operations. 

296 Descriptive Chemistry. 

Gunpowder is a mixture of potassium nitrate, charcoal, and sulphur. 
The ingredients are first purified, pulverized, and thoroughly mixed. 
This mixture is pressed, while damp, into a thin sheet ; and the " press 
cake" thus formed is broken into small grains, which are sorted by 
sieves. The grains are then smoothed or ''glazed" by rolling them in 
a barrel, again sifted, arid finally dried at a low temperature. The pro- 
portions differ with the use of the powder. The United States army 
standard black powder contains 75 per cent of potassium nitrate, 15 of 
charcoal, and 10 of sulphur. When gunpowder burns in a closed space, 
a large volume of gas is suddenly formed. So enormously is this gas 
expanded by the heat that it would fill several hundred times the space 
taken by the powder itself. The pressure exerted by this expanding gas 
is many tons. It is this pressure which forces the ball from a cannon 
and tears a rock to pieces. The chemical changes attending the explo- 
sion of gunpowder in a closed space are complex, as may be seen by the 
following (approximate) equation : 

8 KNO 3 + 90 + 38 = 2 K,C0 3 + K 2 SO 4 -f K 2 S., + 7 CO 2 + 4 N 2 

Probably secondary reactions produce other gases besides carbon diox- 
ide and nitrogen. 

Potassium Chlorate, KC1O 3 , is a white, crystallized, lus- 
trous solid. It tastes like potassium nitrate. It melts at 
334 C., and at a high temperature decomposes into oxygen 
and potassium chloride as final products, thus 

KC1O 3 KC1 + O 3 

Potassium Chlorate Potassium Chloride Oxygen 

It is used to prepare oxygen, and in the manufacture of 
matches and fireworks. In the form of " chlorate of potash 
tablets " it is used as a remedy for sore throat. 

Potassium chlorate is manufactured by passing chlorine into calcium 
hydroxide (milk of lime) and adding potassium chloride to the mixture. 
The simplest equations for the complex changes may be written thus : 

(i) 6 Ca(OH) 2 + 6 C1 2 = Ca(ClO 3 ) 2 + 5 CaCl 2 + 6 H,O 
Calcium Calcium Calcium 

Hydroxide Chlorate Chloride 

Sodium, Potassium, and Lithium. 297 

(2) 3 Ca(ClO) 2 Ca(C10 3 ) 2 + 2 CaCl 2 
Calcium Hypochlorite Calcium Chlorate 

(3) Ca(ClO 3 ) 2 + 2 KC1 2 KC1O 3 + CaCl 2 

Potassium Chlorate 

The salt is also made by the electrolysis of a hot solution of potassium 
chloride, though it has been found more satisfactory to first prepare 
sodium chlorate and convert this salt into potassium chlorate by po- 
tassium chloride. 

Potassium Carbonate, K 2 CO 3 , is a white powder. It 
deliquesces in the air, is very soluble in water, and the 
solution has a strong alkaline reaction. It was formerly 
obtained by treating wood ashes with water, and evaporating 
the solution to dryness. The crude salt thus obtained has 
long been called potash, and a purer product is known 
as pearlash. (The term potash is sometimes applied to 
potassium oxide, K 2 O.) It is used extensively in the manu- 
facture of hard glass, soft soap, caustic potash, and other 
potassium compounds. 

Potassium carbonate is obtained from suint by igniting the greasy 
mass and extracting the potassium carbonate with water. Beet-sugar 
residues also furnish potassium carbonate. After the sugar has been 
obtained from the beet sirup, the molasses is changed by fermentation 
into alcohol, which is distilled off; the liquid residue is evaporated to 
dryness and ignited, and the potassium carbonate extracted with water. 
Pure potassium carbonate is prepared by igniting cream of tartar made 
from the deposits in wine casks. All these sources emphasize the inti- 
mate relation of potassium compounds to vegetable and animal life. 
The bulk of the potassium carbonate is now made from potassium sul- 
phate or from the chloride by the Leblanc process, owing to the abun- 
dance of crude potassium salts at Stassfurt. 

Potassium Hydroxide or Caustic Potash, KOH, is a 

white brittle solid, resembling caustic soda. It absorbs 
water and carbon dioxide very readily ; and if exposed 
to the air, soon becomes a thick solution of potassium 

298 Descriptive Chemistry. 

carbonate. Like sodium hydroxide, it dissolves in water 
with evolution of heat, forming a strongly alkaline caustic 
solution. It is one of the strongest bases, even glass and 
porcelain being corroded by it. Besides its use in the labo- 
ratory, large quantities are consumed in making soft soap. 

Potassium hydroxide is made and purified in the same way as sodium 
hydroxide, viz. by adding lime or milk of lime to a boiling dilute solution 
of potassium carbonate, the equation for the change being : 

Ca(OH) 2 -f K,CO 3 2 KOH + CaCO 3 

Milk of Potassium Potassium Calcium 

Lime Carbonate Hydroxide Carbonate 

It is also made by the electrolysis of a solution of potassium chloride. 

Miscellaneous. Potassium Cyanide (KCN) is a white solid, very 
poisonous, very soluble in water, and having an odor like bitter almonds 
(see Cyanogen, Chapter XIV) . Potassium Sulphate (K 2 SO 4 ) is manu- 
factured from kainite, and is largely used as a fertilizer and in making 
potassium carbonate. 

Relation of Potassium to Life. Potassium, like nitro- 
gen and phosphorus, is essential to the life of plants and 
animals. The ash of many common grains, vegetables, and 
fruits contains potassium as the carbonate. Potassium salts 
are supposed to assist in the formation of starch, just as 
phosphorus is indispensable to the transformation of nitro- 
gen compounds. Potassium salts taken from the soil by 
plants must be returned if the soil is to be productive. 
Sometimes crude kainite is used extensively as a fertilizer ; 
but wood ashes, or the sulphate and chloride, are often 
used to supply potassium salts. 

Lithium, Li, is a silver-white metal and has the specific gravity of 
only 0.59. It is the lightest of the metallic elements. Its compounds 
are widely distributed in small quantities in minerals, mineral waters, 
and plants. Lithia water and citrate of lithium are often prescribed as 
a remedy for diseases of the kidneys. Lithium compounds color the 
Bunsen flame bright red a delicate test for the metal. 

Sodium, Potassium, and Lithium. 299 

Rubidium and Caesium, Rb and Cs, have properties and form com- 
pounds analogous to those of potassium. 


1. Name the alkali metals. What is the symbol of each ? When 
and by whom was each discovered ? 

2. What are the important compounds of sodium ? What per cent 
of the earth's crust is sodium ? 

3. Describe the manufacture of sodium by electrolysis. Describe 
the older methods of manufacture. 

4. Summarize (#) the physical properties, and (<) the chemical 
properties of sodium. How is it usually kept ? For what is it used ? 

5. Discuss the interaction of sodium and water (see Chapter V). 

6. Give the chemical name and formula of common salt Where is 
it found ? 

7. Describe the different methods of preparing salt. State (#) the 
properties, and (b) the uses of salt. 

8. Discuss the manufacture of sodium carbonate by (a) the Le- 
blanc process, (b) By the Solvay process. 

9. What is (a) soda, (b) soda ash, (V) salt cake, (</) soda crystals, 
(V) sal soda, (/*) washing soda, (g) "alkali" ? 

10. State the properties and uses of sodium carbonate. 
n. Describe the preparation, and state (a) the properties, and () 
the uses of sodium bicarbonate. 

12. What is (a) acid sodium carbonate, (b) saleratus, (c} baking 
powder, (d) baking soda, (e) caustic soda ? 

13. State the properties and uses of sodium hydroxide. 

14. Describe the manufacture of sodium hydroxide (a) from lime 
and sodium carbonate, and (b) by electrolysis of sodium chloride. 

15. How is sodium sulphate manufactured? State its properties 
and uses. 

1 6. Where is sodium nitrate found ? State its properties and uses. 

17. Review briefly (a) sodium thiosulphate, (b} water glass, (c) 

1 8. What is a simple test for (a) sodium, and () potassium ? 

19. Give the formula of (#) sodium carbonate, (<) sodium chloride, 
(c) sodium sulphate, (d) sodium hydroxide, (e) sodium bicarbonate, 
(/) Glauber's salt, (g) sodium nitrate. 

300 Descriptive Chemistry. 

20. Discuss the occurrence of potassium compounds. 

21. Discuss the Stassfurt deposits. . 

22. How is potassium prepared ? State (a) its physical properties, 
and () its chemical properties. 

23. Describe the interaction of potassium and water. 

24. Describe the preparation, and state the properties and uses of 
(#) potassium chloride, and (<) potassium nitrate. 

25. Compare potassium nitrate and potassium nitrite. 

26. Describe the manufacture of gunpowder. Upon what does its 
use depend ? 

27. State the properties and uses of (#) potassium chlorate, ($) 
potassium carbonate, (c) potassium hydroxide. 

28. Describe the manufacture of (a) potassium chlorate, () potas- 
sium carbonate, (c) potassium hydroxide. 

29. What is (a) potash, () pearlash, (c} chlorate of potash ? 

30. Discuss the relation of potassium to life. 

31. State the derivation of the names {a) sodium, and (b) potassium. 

32. What is (a) niter, () saltpeter, (c) Chili saltpeter ? 

33. What is the formula of the following compounds of potassium : 
(#) hydroxide, (b) carbonate, (c) nitrate, (</) nitrite, (e) sulphate, (/) 
chlorate, (g) cyanide ? 

34. Describe lithium. For what are its compounds used ? 


1. How much potassium carbonate is necessary to prepare a kilo- 
gram of potassium hydroxide? (Assume K 9 CO 3 + Ca(OH).,= 2 KOH 
+ CaC0 8 .) 

2. What per cent of Glauber's salt, Na 2 SO 4 . ioH 2 O, is sodium 
sulphate ? 

3. A gram of gunpowder produced 300 cc. of gas at o C. What 
would be the volume at 2300 C. ? 

4. How much sodium will 2 kg. of sodium carbonate yield, if heated 
with carbon ? (Assume Na 2 CO 3 + C 2 = Na 2 + 3 CO.) 

5. What is the per cent of sodium in (a) NaOH/ () Na 2 SO 4 , (c} 
NaCl, (d) HNaSO 4 ? 

6. What is the per cent of potassium in (#) potassium bromide 
(KBr), () potassium nitrate (KNO 3 ), (c) potassium iodide (KI) ? 


Introduction. r These metals are related, but they do 
not form a group having such marked family character- 
istics as the alkali metals. The metals, as well as their 
alloys and compounds, have many domestic and commer- 
cial uses. 


Copper has been known for ages. Domestic utensils 
and weapons of war containing copper were used before 
similar objects of "iron. The Romans obtained copper 
from the island of Cyprus/ They called it cuprium aes 
(i.e. Cyprian brass), which finally became simply cuprum. 
From cuprum we obtain the symbol Cu and the terms cu- 
prous and cupric. 

Occurrence of Copper. Copper, both free and com- 
bined, is an abundant element. Single masses of native 
or metallic copper weighing many tons are found in Michi- 
gan mines on the shores of Lake Superior. The most 
valuable ores of copper are copper sulphide (chalcocite, 
copper glance, Cu 2 S), copper oxide (cuprite, ruby ore, 
Cu 2 O), the copper-rron sulphides (copper pyrites, chal- 
copyrite, CuFeS 2 , and bornite, Cu 3 FeS 3 ), and the conir 
plex carbonates (malachite, CuCO 3 Cu(OH) 2 , and azurite, 
2 CuCOg . Cu(OH) 2 ). 

Native copper conies chiefly from Michigan (Fig. 71), the copper- 
iron sulphide ores from Montana, and the carbonates from Arizona. 


302 Descriptive Chemistry. 

The United States produced about 300,000 tons of copper in 1902, 
which was more than half of the world's supply. Of this amount Mon- 
tana furnished about 38 per cent, Michigan 26 per cent, and Arizona 22 
per cent. The annual output has steadily increased since 1896. 

Metallurgy of Copper. Copper is extracted from its 
ores by processes which vary with the composition of the 
ore. (i) Native copper ore is first crushed, then washed 
to remove impurities, and the concentrated product finally 
smelted and refined by a single fusion. (2) The carbon- 
ates and oxides are reduced by roasting them with coke in 
blast furnaces. The general chemical change may be rep- 
resented thus 

Cu 2 O + C = 2Cu + CO 

Copper Oxide Carbon Copper Carbon Monoxide 

(3) The smelting of copper-iron sulphides is complicated. 
The ore is crushed and washed, and then roasted in a fur- 
nace. This operation removes the adhering rock and 
changes much of the sulphide into an oxide. The roasted 
mass is then melted with coal and sand in a shaft or a 
reverberatory furnace, whereby the iron is largely changed 
into a fusible silicate, which runs off as a part of the slag. 
The remaining "matte," as it is called, contains from 50 
to 65 per cent of copper, besides some iron, sulphur, and 
arsenic. It is roasted and melted until all the iron and 
arsenic are removed and mainly copper sulphide remains. 
This is finally roasted to convert it partly into an oxide, 
and the mixture of sulphide and oxide is again melted ; the 
sulphur passes off as sulphur dioxide, and the copper is 
left behind. The equation for this final change is 

2 CuO + Cu 2 S = 4 Cu -h SO 2 
Copper Oxide Copper Sulphide Copper Sulphur Dioxide 

Copper Silver Gold. 































Sometimes the sulphur and arsenic are removed by forcing 
hot air through the molten sulphide. 

Purification of Copper. The crude copper from most 
ores contains about 98 per cent of copper. Such impure 
copper is best purified by electrolysis, and is called electro- 
lytic copper. Thick plates of the impure copper are 
attached as anodes to the 
positive electrode of a 
powerful battery or dy- 
namo and suspended in 
a solution of copper sul- 
phate and sulphuric acid. 
Sheets of pure copper are 

attached as Cathodes tO FIG. 57. Apparatus for the preparation 

the negative electrode and, *& * Shodet * * *" 
of course, dip into the 

solution (Fig. 57). When the current passes, the crude 
copper anodes dissolve, pure copper is deposited upon the 
cathodes, and the impurities either remain in solution or 
fall to the bottom of the tank as mud. From this mud, 
gold and silver are extracted in appreciable quantities. 
Electrolytic copper is very pure. 

Properties of Copper. Copper is a bright metal, dis- 
tinguishable from all others by its peculiar reddish color. 
It is flexible, hard, and tough ; it can be drawn out into 
wire and rolled into very thin sheets. Its specific grav- 
ity is 8.9. Next to silver, copper is the best conductor of 
heat and electricity. Exposed to dry air, it turns dull, and 
in moist air it gradually becomes coated with a greenish 
copper carbonate. Heated in the air, it is changed into 
the black copper oxide, and at a high temperature it colors 
a flame emerald green. With nitric acid it forms copper 

304 Descriptive Chemistry. 

nitrate and oxides of nitrogen (see Oxides of Nitrogen); 
with hot sulphuric acid it yields copper sulphate and sul- 
phur dioxide (see Sulphur Dioxide). Hydrochloric acid 
has little effect upon it. Copper replaces some metals if 
suspended in solutions of their compounds, e.g. a clean 
copper wire soon becomes coated with mercury if placed 
in a solution of any mercury compound ; on the other 
hand, metals like iron, zinc, and magnesium remove cop- 
per from its solution, e.g. a nail or knife blade soon becomes 
coated with copper if dipped into a solution of any copper 
compound. Scrap iron is often used to precipitate copper 
on a large scale. 

Test for Copper. (i) The reddish color, peculiar "coppery" taste, 
and green color given to a flame serve to identify metallic copper. 
(2) An excess of ammonium hydroxide added tothe solution of a cop- 
per compound produces a beautiful blue solution. (3) A few drops 
of acetic acid and of potassium ferrocyanide solution added to a dilute 
solution of a copper compound produce a brown precipitate of copper 
ferrocyanide. These tests are characteristic and decisive. 

Uses of Copper. Next to iron, copper is the most use- 
ful metal. Enormous quantities of wire are used in operat- 
ing the telegraph, cable, telephone, electric railway, and 
electric light. Sheet copper is made into household utensils 
boilers, and stills. Copper bolts, nails, and rivets are used 
on ships, because the rust does not destroy wood as iron rust 
does. All nations use copper as the chief ingredient of 
small coins. Electrical apparatus utilizes much copper. 
Maps, etchings, and some kinds of engravings are printed 
from copper plates ; calico is printed from a copper cyl- 
inder upon which the design is engraved. Books are 
printed and illustrated from an electrotype, made by de- 
positing a film of copper upon an impression of the type 
or design in wax or plaster of Paris. In a similar way 

Copper Silver Gold. 305 

many objects are copper plated (see Chapter X). Copper 
is an essential constituent of many alloys. 

Alloys of Copper are important. Brass is a bright yel- 
low alloy containing 63 to 72 per cent of copper, the re- 
mainder being zinc. It is made by melting these metals 
together. It can be drawn into wire, hammered into any 
shape, and turned in a lathe. It is harder than copper, 
and on account of its durability and elasticity has many uses 
for which copper is not suited. Pinchbeck, Muntz metal, 
Bath metal, Dutch metal (leaf or "gold"), are varieties of 
brass. Muntz metal is now used in place of sheet copper, 
as sheathing for the bottoms of ships, because it rusts very 
slowly. Typical bronze contains different proportions of 
copper, zinc, and tin. Some antique bronzes contain lead 
or iron. The per cent of copper is 70 to 95, of zinc I to 
25, of tin i to 1 8. The proportions in the British bronze 
coinage are copper 95, zinc i, tin 4. On account of its 
beautiful color and extreme durability, bronze is used for 
statues, memorial tablets, coins, and medals. The ancients 
made it into weapons of war and household utensils. Can- 
non were formerly made of bronze, but for this purpose 
steel is now used. Phosphor bronze contains a small per 
cent of phosphorus and of lead. It is tougher than ordi- 
nary bronze, and is used to make steamship propellers and 
parts of machines. Silicon bronze is copper with traces of 
iron and silicon, and is used for telegraph and telephone 
wires. Aluminium bronze contains 90 per cent copper 
and 10 per cent aluminium. It is a hard, yellow, elastic 
alloy, and is used in constructing hulls of yachts ; its light- 
ness, strength, and resistance to chemicals adapt it to many 
other uses. 

Gun metal is about 90 per cent copper and 10 per cent zinc ; it was 
formerly used in making cannon, and is now used to some extent in 

306 Descriptive Chemistry. 

making firearms. Bell metal contains about 75 per cent copper and 
25 per cent zinc. Speculum metal contains about 70 per cent copper, 
30 per cent tin, and traces of zinc, nickel, and iron ; it takes a brilliant 
polish, and is used in optical instruments, especially telescopes, to re- 
flect light. The different varieties of German silver contain different 
proportions of copper, nickel, and zinc. The per cent of copper is 50 to 
60, of nickel 20 to 25, and of zinc about 20. In color and luster it re- 
sembles silver, for which it is often substituted. Its power to conduct 
electricity is only slightly affected by changes of temperature, hence 
it is often used in resistance coils. Chinese Pakfong (or paktong) is a 
variety of German silver. The nickel coins of Germany and the 
United States contain 75 per cent copper and 25 per cent nickel. Cop- 
per is also a constituent of many other coins. Britannia metal and 
white metal, in which copper is a minor constituent, are described 
under Alloys of Tin. 

Compounds of Copper. Copper forms two series of com- 
pounds, the cuprous and the cupric. Thus, there are 
cuprous oxide (Cu 2 O) and cupric oxide (CuO), cuprous 
chloride (CuCl) and cupric chloride (CuCl 2 ). The cuprous 
compounds contain a larger proportion of copper than the 
cupric compounds. Not every member of each series is 
important, or even well known. Other metals mercury 
and iron form similar series. The most important com- 
pounds are the oxides and copper sulphate. Copper com- 
pounds are poisonous. Cooking utensils made of copper 
should be used with care. Vegetables, acid fruits, and 
preserves, if boiled in them, should be removed as soon as 
cooked. The vessels themselves should be kept bright, to 
prevent the formation of soluble copper salts, which might 
contaminate the contents. 

Cuprous Oxide, Cu 2 O, occurs native as cuprite or ruby ore. It may 
be obtained as reddish powder by heating a mixture of solutions of cop- 
per sulphate, Rochelle salt, sodium hydroxide, and grape sugar. This 
oxide colors glass ruby red. It is a beautiful mineral and a valuable 

Copper Silver Gold. 307 

Cupric Oxide, CuO, is a black solid formed by heating 
copper nitrate. It is reduced to metallic copper by hydro- 
gen or by carbon, thus 

CuO + H 2 = Cu + H 2 O 
Cupric Oxide Hydrogen Copper Water 

Hence it may be used to determine the gravimetric com- 
position of water. 

Copper Sulphate, CuSO 4 , is the most useful compound 
of copper. Like many of the cupric compounds it is a blue, 
crystallized solid, and is often called " blue vitriol " or 
" blue stone." The crystallized salt (CuSO 4 . 5 H 2 O) loses 
water in the air ; heated to 240 C, all the water escapes, 
leaving a whitish powder. This anhydrous copper sul- 
phate absorbs water from alcohol and similar liquids^^nd 
when added to water it again becomes blue. Copper -sul- 
phate is used in electric batteries, in making other copper 
salts, in calico printing, dyeing, copper plating, in preserv- 
ing timber, and whenever a soluble copper compound is 
needed. It is poisonous and is one ingredient of certain 
mixtures which are sprayed upon trees to kill insects. 

Copper sulphate may be prepared by treating copper 
with sulphuric acid. This method is used on a large scale, 
but much of the copper sulphate of commerce is a by- 
product obtained in refining gold and silver with sulphuric 
acid (see below). 

Copper Nitrate, Cu(NO 3 ) 2 , is a blue, crystallized solid, formed by the 
interaction of copper and dilute nitric acid. It is a cupric salt. It is 
very soluble in water, and is readily decomposed by heat into cupric 
oxide and oxides of nitrogen. 

Cuprous Sulphide, Cu 2 S, is the bluish black mineral chalcocite. Cu- 
pric sulphide, CuS, is the black precipitate formed by passing hydrogen 
sulphide gas into a solution of a cupric salt. 

308 Descriptive Chemistry. 

Malachite is a bright green mineral and is often used as an orna- 
mental stone. Azurite is a magnificent blue, crystallized mineral. 
Both are carbonates and valuable ores of copper. 


Silver is one of the precious metals. From the remotest 
ages it has been used for ornaments, household vessels, 
and money. 

The Latin name of silver is argentine, from which the symbol Ag is 
derived. The alchemists called it luna, on account of its silvery or 
" moonlike" appearance. 

Occurrence of Silver. Native silver is found in Ari- 
zona, Mexico, Norway; also in South America and Aus- 
tralia. The chief ores are the sulphides. The simple 
sulphide (silver glance, argentite, Ag 2 S) is the richest ore 
and is found in many localities in the United States. Sil- 
ver sulphide is often combined with sulphides of lead, 
copper, antimony, or arsenic. These complex sulphides 
are found in Mexico, Peru, Bolivia, Chili, and in Idaho. 
Small quantities of native silver chloride (horn silver, 
AgCl) are also found ; it resembles wax or horn, and melts 
in a candle flame. Sea water contains traces of silver, the 
total quantity in the ocean being estimated to be about two 
million tons. Alloys of silver with gold, mercury, and 
copper are found ; average California gold contains about 
12 per cent silver. Many ores contain silver, especially 
those of lead; and this argentiferous (or silver-bearing) 
lead is one of the chief sources of silver. 

The world's supply of silver comes mainly from the United States. 
Mexico, Germany, Australia, and Bolivia. The United States produced 
over sixty-four million ounces in 1902. This was about one third of the 
world's supply, and also the average annual output for the last few 
years. Of this vast quantity, about 90 per cent was furnished by Colo- 
rado, Montana, Utah, Idaho, California, and Nevada (Fig. 58). 

Copper Silver Gold. 


Metallurgy of Silver. Silver is extracted from its 
ores by two principal processes, (i) In the amalgama- 
tion process the powdered ore is first changed into silver 

FIG. 58. Distribution of silver and gold in the United States. 

chloride by roasting (or simply mixing) it with sodium 
chloride. The mass is then reduced to silver by agitation 
with water and iron (or an iron compound) ; the simplest 
equation for this reaction is 

2 AgCl + 
Silver Chloride 

Fe = 



FeCl 2 

Iron Chloride 

The silver is removed by adding mercury, which forms an 
amalgam (an alloy) with the silver, but not with the other 
substances. When the amalgam is heated, the mercury 
distils off, and the silver with some gold remains be- 
hind. (2) Silver is extracted from lead ores by the Parkes 
process. After the sulphur, arsenic, and other impurities 
have been removed from the lead ores, the final product is 
a mixture of lead, silver, and gold. This is melted and 

310 Descriptive Chemistry. 

thoroughly mixed with zinc. As the mixture cools, an 
alloy of silver, gold, zinc, and a little lead rises to the top, 
solidifies, and is removed. The remaining lead mixture is 
treated again with zinc. The alloy of silver, gold, zinc, 
and lead is heated to volatilize the zinc and to oxidize (or 
melt away) the lead. The mixture of silver and gold is 
heated with sulphuric acid ; the gold is not acted upon, 
but the silver forms silver sulphate, which is reduced by 
copper to metallic silver (Fig. 59). 

199-7 7 

FIG. 59. Bar or " brick " of silver showing the stamp of the United States Assay 
Office as a guarantee of its purity. 

o r 
Office as a guarantee of its purity 

Lead ores containing considerable silver are sometimes subjected to 
CUpellation to extract the silver. The ore or alloy is heated in a fur- 
nace having a shallow hearth made of porous, infusible bone ash. The 
lead is changed into an oxide (litharge), which melts, and is partly 
driven off by the air blast into pots and partly absorbed by the porous 
cupel. The silver is protected from the oxidizing power of the air by 
the melted litharge, but toward the end of the operation the thin film of 
litharge bursts, and the metallic silver appears as a bright disk if the 
operation is conducted in a furnace, and as a globule or button if the 
extraction is performed in a small assay cupel. The process is then 
stopped and the silver removed. 

Properties of Silver. Silver is a lustrous, white metal, 
which takes a brilliant polish. It is harder than gold, but 
softer than copper. Like copper, it is ductile and malle- 
able, and may be easily made into various shapes. Its 
specific gravity is about 10.5, being heavier than copper, 

Copper Silver Gold. 311 

but lighter than lead. It melts at about 954 C, and fuses 
readily on charcoal in the blowpipe flame ; it vaporizes in 
the oxyhydrogen flame and in the electric furnace. Molten 
silver absorbs about twenty times its volume of oxygen, 
which is expelled violently when the silver solidifies. Pure 
silver conducts heat and electricity better than any other 
metal, but it is too costly for such uses. It does not tarnish 
in air, unless sulphur compounds are present, and then the 
familiar black film of silver sulphide is produced. This 
blackening is especially noticed on silver spoons which 
have been put into eggs or mustard, and on silver coins 
which have been carried in the pocket, the sulphur in the 
latter case coming from sulphur compounds in the perspira- 
tion ; the tarnishing of household silver is due to sulphur 
compounds in illuminating gas or gas from burning coal. 
So-called " oxidized " silver is not oxidized, but coated with 
silver sulphide. Silver is only very slightly acted upon by 
hydrochloric acid, and not at all by molten caustic potash, 
soda, or potassium nitrate. Nitric acid and hot concen- 
trated sulphuric acid change it into the nitrate and sulphate, 
respectively, as in the case of copper. 

Alloys of Silver. Pure silver is too soft for constant 
use, and is usually hardened by adding a small amount of 
copper. These alloys are used as coins and for jewelry. 
The silver coins of the United States and France contain 
900 parts of silver to 100 of copper, and are called 900 
fine. British silver coins are 925 fine ; this quality is called 
"sterling silver," and from it much ornamental and useful 
silverware is made. 

Silver Plating. Metals cheaper than silver may be 
coated or plated with pure silver precisely as in the case of 
copper. Plated silverware has the appearance of solid or 
pure silver. The object to be plated is carefully cleaned, 

312 Descriptive Chemistry. 

and made the cathode in a bath or solution of potassium 
silver cyanide. The anode is a plate of pure silver (Fig. 

60). The deposit of silver is 
dull, but may be brightened 
by rubbing with or without 

Compounds of Silver. - 
FIG. 60. Apparatus for silver plat- The most important compound 

ing. A, A, A, are silver anodes, and . ., . , / . AT ^ x 

thf spoons are cathodes. SllvCI nitrate (AgNO 3 ). 

It is a white crystalline solid, 

made by dissolving silver in nitric acid. Exposed to the 
light, it turns dark if in contact with organic matter. It 
discolors the skin ; if applied long enough, it disintegrates 
the flesh, and is often used by physicians for this purpose. 
Its caustic action and the silvery color of the metal from 
which it is made long ago led to its name, lunar caustic. 
Besides its extensive use in photography and silver plating, 
silver nitrate is the essential constituent of indelible ink. 
Silver chloride (AgCl) is made by adding hydrochloric 
acid or the solution of any chloride to a solution of a silver 
compound. Thus formed, it is a white, curdy solid, which 
turns violet in the light, and finally black. This action of 
light is more intense if organic ntatter is present. It 
dissolves in ammonium hydroxide, forming a complex com- 
pound of the two substances. The formation and proper- 
ties of silver chloride constitute the test for silver. Silver 
bromide (AgBr) and silver iodide (Agl) are analogous to 
silver chloride in their properties and methods of forma- 
tion. They are used in photography. 

Photography is based on the fact that silver salts, espe- 
cially the bromide and iodide, change color when mixed 
with organic matter and exposed to the light. The photo- 
graph is taken on a glass plate, coated on one side with a 

Copper Silver Gold. 313 

thin layer of gelatine, containing the silver salts. Some- 
times a sheet of sensitized gelatine, called a film, is used. 
The plate or film is placed in the 'camera and exposed. 
The light, which comes from the object being photographed, 
changes the silver salts in proportion to its brilliancy. The 
plate, however, shows no change until it has been devel- 
oped. This process consists in treating the plate with a 
reducing agent, e.g. ferrous sulphate, pyrogallic acid, or 
special mixtures. As the developer acts upon the plate, 
the image appears. This is really a deposit of finely 
divided silver. Where the intense light fell upon the 
plate, the deposit is heavier than where little or no light 
fell. Hence, dark parts of the object appear light on the 
plate, and light parts dark; and since the image is the 
reverse of the object, the plate is called a negative. When 
the plate has been properly developed, it still contains sil- 
ver salts not altered by the light ; and if they were left on the 
plate, the image would be clouded, and finally obliterated 
by the light. The image is, therefore, fixed by wash- 
ing off the silver salts with a solution of sodium thiosul- 
phate (or "hyposulphite"). A print is made by laying 
sensitized paper upon the negative and exposing them to 
the sunlight, so that the light will pass through the nega- 
tive. The negative obstructs the light in proportion to the 
thickness of the silver deposit, so the photograph has the 
same shading as the object. Most prints, like the plates, 
must be fixed. Sometimes the color is improved by toning, 
i.e. by placing the print in a solution of gold or of platinum. 


Gold is the most precious of the metals, and has 
been used from the earliest times for adornment and as 

314 Descriptive Chemistry. 

The Latin name of gold, aurum, gives the symbol Au. For many 
centuries the alchemists tried to produce gold from base or cheaper 
metals. They were unsuccessful in their search for the Philosopher's 
Stone, which they believed had power to affect this transformation. 

Occurrence of Gold. Gold is widely distributed, but 
not abundantly in many places. Unlike copper and silver, 
its compounds are few and rare; the only important ones 
are the tellurides (compounds of tellurium) found in Colo- 
rado. It is never found pure, being alloyed with silver 
and occasionally with copper or iron. It is disseminated 
in fine, almost invisible, particles among ores of other 
metals, though not so abundantly as silver. Much gold is 
found in veins of quartz, and in the sand and gravel formed 
from gold-bearing rocks. Gold occurs usually as dust, 
scales, or grains, but occasionally shapeless masses called 
" nuggets " are found, varying in weight from a few grams 
to many kilograms. The largest nugget ever known 
weighed over 84 kg. (184 Ibs.). 

The chief gold-producing countries are the United States, Australia, 
South Africa, and Russia. In 1902 the United States produced nearly 
four million ounces, which came largely from Colorado, California, and 
other Western states, and Alaska. Gold in working quantities is found 
in about twenty states of the Union (Fig. 58). The total value of the 
gold produced in the world in 1902 was about $306,000,000. 

Gold Mining. Gold was first obtained by miners by 
washing the gold-bearing sand and gravel of a stream in 
large pans or cradles. This primitive method was soon re- 
placed by placer mining and hydraulic mining. Streams 
of water, directed against the earth containing the gold, 
wash away the lighter materials, but leave the heavy gold 
behind in the form of scales or " gold dust." From this 
mixture gold and silver are extracted by mixing with mer- 
cury, or by passing the moistened mass over copper plates 

Copper Silver Gold. 315 

coated with mercury. The amalgam is then heated, as in 
the metallurgy of silver, to remove the mercury ; the resi- 
due of gold and silver is purified as described below. In 
vein mining the gold-bearing rock usually quartz is 
crushed and then washed, and the gold removed by mer- 
cury, as in placer mining (see Chapter XX). Low grade 
ores and those containing certain metals cannot be profita- 
bly treated with mercury. In the chlorination process the 
crushed ore is roasted and then revolved in barrels contain- 
ing bleaching powder and sulphuric acid; this operation 
forms a soluble gold chloride ( AuCl 3 ), from which the gold 
is precipitated as a fine powder by hydrogen sulphide (or 
other reducing agents). In the cyanide "process the 
crushed ore, or the slime from a previous extraction, is 
mixed with a weak solution of potassium cyanide and 
exposed to the air ; this operation changes the gold into a 
soluble cyanide (KAu(CN) 2 ). The gold is separated from 
this solution by electrolysis or by treatment with zinc. 

Purification of Gold. Gold obtained by the above 
methods is impure, silver being the chief impurity. These 
metals are parted by a chemical process or separated by 
electrolysis. By the old parting process known as quar- 
tation an alloy of gold and silver, in which the gold is 
about one fourth of the whole, is treated with nitric acid ; 
this operation changes the silver into the nitrate from 
which the pure gold may be readily removed. The metals 
may be parted by the cheaper method described under 
silver, viz. by boiling with concentrated sulphuric acid. 
By this treatment the gold, which is about one sixth of the 
alloy, is left as a brownish, porous mass. It is washed, 
dried, and fused with charcoal and sodium carbonate. In 
the electrolytic method of separation, the anode is an 
alloy of gold and silver, the cathode is silver, and the elec- 

3 1 6 Descriptive Chemistry. 

trolyte is nitric acid. When the current passes, part of the 
silver of the anode goes into solution as the nitrate, while 
part is deposited at the cathode ; the gold remains at the 
anode as a fine powder and is caught in a cloth bag which 
incloses the whole anode. Gold is now purified at the 
United States Mint by electrolysis. The electrolyte is a 
hydrochloric acid solution of gold chloride, but otherwise 
the process is the same as described above. It is more 
economical than parting by nitric acid. 

The purity of gold is expressed in carats. Pure gold is 
24 carats fine ; an alloy containing 22 parts of gold and 2 
parts copper is 22 carat gold, while one containing equal 
parts gold and other metals is 12 carat gold (see foot-note, 
page 183). 

Properties of Gold. Gold is a yellow metal. It is 
about as soft as lead, and is the most ductile and malleable 
of all metals. The leaf into which it may be beaten is 
very thin and is green by transmitted light. Air, oxygen, 
and most acids do not attack it ; but it is changed into a 
gold chloride (AuCl 3 ) by aqua regia (see Aqua Regia). 
Gold is one of the heaviest metals, its specific gravity 
being about 19. 

Uses of Gold. Pure gold is too soft for most practical 
purposes, and is, therefore, usually hardened with copper 
or silver. The gold-copper alloy has a reddish color and 
is often called " red gold " ; the gold-silver alloy is paler 
than pure gold and is sometimes called "white gold." 
Gold coins contain gold and copper. The United States 
standard gold coins contain 9 parts gold and I part cop- 
per, while in England the legal standard is n of gold to 
i of copper. Gold leaf of various grades is used to orna- 
ment books, signs, and many objects. Jewelers use gold 
for many purposes; such gold varies from 1 2 to 22 carats 

Copper Silver Gold. 317 

in purity. On account of its malleability, feeble chemical 
action, and beauty, gold is used by dentists for filling 

Compounds of Gold are readily decomposed by metals, weak reducing 
agents (e.g. ferrous sulphate or hydrogen sulphide), fine solids like char- 
coal, and by electrolysis. When gold is dissolved in aqua regia and 
the acid removed by evaporation, the resulting gold chloride (AuCl 3 ) 
gives with stannous chloride solution a beautiful purple precipitate ; the 
latter is called "purple of Cassius," and is probably finely divided gold. 
Its formation is the test for gold. The process of gold plating is the 
same as silver plating, only the solution is one of potassium gold cyan- 
ide (Au(CN) 3 . KCN) and the anode is gold. Much cheap jewelry is 
gold plated. 


1. What is the symbol of (a} copper, (b) silver, (c} gold? State 
the derivation of each symbol. 

2. Where is copper found abundantly? State in what form it 
occurs in each locality. Discuss its production. 

3. Describe briefly the metallurgy of (a) native copper, () oxides 
and carbonates, (c) copper-iron sulphides. 

4. Describe the purification of copper by electrolysis. 

5. State (#) the physical properties of copper, and () the 
chemical properties. 

6. Describe several tests for copper. 

7. Discuss the uses of copper. 

8. Name ten alloys of copper. Describe five important alloys. 

9. What is an electrotype? How is it made? (See Chapter X.) 

10. State the general properties of copper compounds. 

11. Describe the oxides of copper. 

12. Describe the manufacture, and state the properties and uses of 
copper sulphate. 

13. What are the properties of (a) copper nitrate, (b} malachite, 
and (<:) azurite? 

14. What is the formula of (#) copper sulphate, ($) copper nitrate, 
(c) cupric oxide, and (*/) cuprous oxide? / 

15. Discuss (#) the occurrence, and (<) the production of silver. 

1 6. Describe the extraction of silver by (a) the amalgamation pro- 
cess, and () the Parkes process. 

318 Descriptive Chemistry. 

17. State (a) the physical properties of silver, and () the chemical 

1 8. Discuss (#) silver alloys, and (b) silver plating. 

19. State the properties and uses of silver nitrate. 

20. State the properties of silver chloride. What is the test for 

21. Describe briefly the essential operations in photography. What 
general chemical changes does it utilize? 

22. What is (a} blue vitriol. (<) argentiferous lead, (c) oxidized 
silver, (d) sterling silver, (e) coin silver, (/) lunar caustic, and (g) 

23. What is the formula of (a) silver nitrate, and () silver chloride? 

24. Discuss (a) the occurrence of gold, and ($) its production. 

25. Describe the different methods of (a) mining, and () extracting 

26. Describe the purification of gold by (a) parting, and () elec- 

27. What is 1 8 carat gold? 

28. State (a} the properties, and (/) the uses of gold. 

29. Discuss (#) compounds of gold, and (b} gold plating. 

30. What is the test for gold? 

31. What is (a) gold dust, (<) aqua regta, (V) a nugget, (d) gold 


1. How much cupric oxide is formed by heating 1467 gm. of copper 
in air? (Assume Cu + O = CuO.) 

2. Calculate the per cent of copper in (a) malachite (CuCO . 
Cu(OH) 9 ), () azurite (2 CuCO 3 . Cu(OH) 9 ), (c) copper sulphate 
(CuS0 4 ). 

3. If 480 gm. of silver interact with nitric acid, how much silver 
nitrate is formed? 

4. Calculate the per cent of silver in (a} silver chloride (AgCi), 
() silver sulphide (AgS 2 ), (c) silver nitrate (AgNO 3 ). 


THESE elements form a natural group called the alkaline 
earth metals. The metals themselves are rare, but their 
compounds, especially those of calcium, are numerous and 
useful. This group resembles the alkali group. 


Occurrence of Calcium. Calcium is never found free. 
Combined calcium makes up about 3.5 per cent of the 
earth's crust. The most abundant compound is calcium 
carbonate (CaCO 3 ). This has many familiar forms, e.g. 
limestone, chalk, marble, coral, and shells. Many rocks 
are complex silicates of calcium and other metals. The 
extensive deposits of calcium phosphate, calcium borate, 
and calcium fluoride have been mentioned. Calcium sul- 
phate (CaSO 4 ) occurs abundantly in the form of gypsum, 
alabaster, and selenite. Calcium compounds are essential 
to the life of plants and animals, being found in the leaves 
of plants, and in the bones, teeth, and shells of animals. 
Many rivers and springs contain calcium salts, especially 
the acid carbonate and sulphate. 

Preparation and Properties. Metallic calcium was obtained by 
electrolysis in 1808 by Davy, but our knowledge of the pure metal is 
due to Moissan. In 1898 he prepared it from the iodide by electrolysis 
and by fusion with sodium. The equation for the latter process is 

CaI 2 + 2 Na = Ca + 2 Nal 

Calcium Iodide Sodium Calcium Sodium Iodide 


Descriptive Chemistry. 

Calcium is a silver-white metal, soft enough to be cut with a knife, 
though harder than potassium. It may be crystallized from melted 
sodium. It readily decomposes water at the ordinary temperature, and 
combines directly with most of the other elements. 

Calcium Carbonate, CaCO 3 . The most abundant form 
of this compound is limestone. Vast deposits are found 
in many places, exhibiting a variety of textures and colors. 
In the United States much limestone is found in Iowa, 
Illinois, and Wisconsin. All kinds are compact and usually 
soft, though some are hard enough for use as building 
stone; some are coarse, and often consist of grains, crys- 
tals, or small shells. Pure limestone is white or gray, but 

impurities, especially organic 
matter and iron compounds, pro- 
duce blue, yellow, reddish, and 
black varieties. Hard, crystal- 
line limestone which takes a good 
polish is called marble. This 
form, which has a wide range of 
color, is used as a building and 
an ornamental stone. Calcite 

FIG. 61. Calcite crystals. 

is crystallized calcium carbonate. It is almost as abundant 

as quartz, though 

softer ; its varied 

color and crystal form 

combine to make it 

attractive (Fig. 61). 

A very transparent 

variety of calcite 

called Iceland spar 

FlG. 62. Crystallized Iceland spar showing 
double refraction. 

has the remarkable 
property of double 
refraction, i.e. of making objects appear double (Fig. 62). 

Calcium, Strontium, and Barium. 321 

Calcium carbonate is not soluble in water, unless carbon 
dioxide is present (see Carbon Dioxide). As water con- 
taining carbon dioxide works its way underground in 
limestone regions the limestone is dissolved and caves 

FIG. 63. Stalactites and stalagmites in Luray Cavern. 
From a photograph copyrighted by C. H. James. 

are often formed or enlarged. When the water enters a 
cave and drips from the top, the water evaporates, or 
the gas escapes, or both, and the calcium carbonate is 
redeposited, often forming stalactites and stalagmites 
(Fig. 63). The stalactites hang from the roof like icicles, 

322 Descriptive Chemistry. 

while the stalagmites grow up from the floor, as the 
deposit slowly accumulates from the solution which 
drops from the roof or the tips of stalactites. The 
Mammoth Cave in Kentucky, the Marengo Cave in Indi- 
ana, and the Luray Cavern in Virginia are famous for 
these fantastic formations. Mexican onyx is a variety of 
stalagmite. Vast deposits of this beautiful mineral are 
found in Algeria and Mexico. It is translucent and deli- 
cately colored, and is used as an ornamental stone, espe- 
cially for altars, table tops, mantels, and lamp standards. 
Beautiful deposits of limestone are found around many 
mineral springs. Travertine occurs near many springs 
in Italy. When fresh, it is soft and porous, but it soon 
hardens and becomes a durable building stone in dry cli- 
mates. The outer walls of the Colosseum and of St. 
Peter's are travertine. Limestone often contains shells 
and fossils, confirming our belief that limestone is the 
remains largely of the shells of animals. The calcium 
carbonate dissolved in the ocean is transformed by marine 
organisms into shells and bony skeletons. The hard parts 
of these animals accumulate in vast quantities on the ocean 
bottom, become compact, often hardened and crystallized, 
and are finally elevated into their present position. On 
the coast of Florida, coquina or shell rock is found. It is 
a mass of fragments of shells cemented by calcium carbon- 
ate, and in time will become compact limestone. Chalk 
is the remains of shells of minute animals. When exam- 
ined under a microscope, a good specimen is seen to con- 
sist almost entirely of tiny shells. The ocean contains 
myriads of minute animals, and when they die, their shells, 
which are calcium carbonate, sink to the bottom. As a 
result, the ocean bottom is partly covered with a gray 
mud, called globigerina ooze. Under the microscope this 

Calcium, Strontium, and Barium. 


ooze looks like Figure 64, and when dried and compressed 
it can hardly be distinguished from chalk. Hence it is be- 
lieved that the immense beds of chalk found in England 
and other places were formed from this ooze. Some vari- 
eties of chalk under the microscope resemble the ooze 

FIG. 64. Ooze from the ocean 
bottom, showing globigerina shells 

FIG. 65. Chalk from Iowa, showing 
globigerina shells (magnified). 

(Fig. 65). Blackboard crayon is a mixture of chalk and 
clay. Whiting is a variety of impure chalk ; putty is a 
mixture of whiting and oil. Coral is calcium carbonate. 
The vast accumulations in the sea are the skeletons of the 
coral animals. 

The properties of calcium carbonate, discussed in Chap- 
ter XIV, may be profitably reviewed at this point. 

Besides being burned into lime, immense quantities of 
limestone are consumed in manufacturing iron and steel, 
the United States alone using annually over seven million 
tons in this industry. 

324 Descriptive Chemistry. 

Calcium Oxide, CaO, is the chemical name of lime. It 
is a hard, white solid. Pure lime is almost infusible, and 
when heated in the oxyhydrogen flame, it gives an in- 
tensely bright light, sometimes called the "lime light" (see 
Hydrogen). In the electric furnace it melts and volatil- 
izes, if the heating is prolonged. Lime containing impuri- 
ties, like sand, clay, and iron compounds, melts quite readily 
into a glass or slag. Exposed to the air, lime becomes 
" air slaked," i.e. it slowly absorbs water and carbon diox- 
ide, swells, and soon crumbles to a powder, which is a 
mixture of calcium hydroxide and calcium carbonate. 
Lime and water combine violently and liberate consider- 
able heat, as is often seen when mortar is being prepared. 
This operation is called " slaking," and the product is 
"slaked lime." The equation for the chemical change 

CaO + H 2 O = Ca(OH) 2 
Calcium Oxide Water Calcium Hydroxide 

Fresh lime attacks organic matter, and is therefore often 
called " caustic lime " or quicklime. It combines with 
water to form calcium hydroxide and with acids to form 
calcium salts. 

Lime is one of the most important substances. It 
is used in preparing mortar, cement, metals, in making 
bleaching powder, calcium carbide, sodium hydroxide, 
and glass, in purifying illuminating gas and sugar, to 
remove hair from hides before the process of tanning, in 
dyeing and bleaching cotton cloth, in drying gases, and as 
a disinfectant and fertilizer. 

Lime is prepared on a large scale by heating limestone 
in a partly closed cavity or vessel. The decomposition 
takes place according to the equation 

Calcium, Strontium, and Barium. 


FIG. 66. Limekiln (ver- 
tical section). The fire is 
built in A under the arch of 

CaCO 3 = CaO + CO 2 
Calcium Carbonate Calcium Oxide Carbon Dioxide 

The carbon dioxide gas escapes and the lime is left in the 

Limestone was formerly "burned" in a 
cavity on a hillside, and in some regions it is 
so prepared to-day. An arch of limestone is 
built across the cavity above the fire pit, and 
limestone is piled upon the arch until the 
kiln is full (Fig. 66). The fire is then lighted 
and kept burning for about three days. These 
kilns have been largely replaced by a modern 
furnace, constructed so that the heat can be 
regulated, the gases swept out, and the prod- 
uct removed without extinguishing the fire. 

Limestone, containing more than 10 per 
cent of clay, forms hydraulic lime, which 

becomes very hard when wet or kept in contact with water. Cements 
are varieties of hydraulic lime. They are made by burning a mixture 
(natural or artificial) of limestone, clay, and sand, and grinding the 
product to a very fine powder. Rosendale and Portland are the 
common brands. The hardening of cements is not well understood. 

Calcium Hydroxide, Ca(OH) 2 , is a white powder. It 
is sparingly soluble in water, but more soluble in cold than 
in warm water. The solution has a bitter taste, an alkaline 
reaction, and is commonly called limewater. Exposed to 
the air, limewater becomes covered with a thin crust of cal- 
cium carbonate, owing to the absorption of carbon dioxide. 
For the same reason, limewater becomes milky or cloudy 
when carbon dioxide is passed into it. The formation of 
calcium carbonate in this way is the usual test for carbon 
dioxide. The equation for this chemical change is 

Ca(OH) 2 + C0 2 = CaC0 3 + H 2 O 
Limewater Carbon Calcium 
Dioxide Carbonate 

326 Descriptive Chemistry. 

Limewater is prepared by carefully adding lime to consid- 
erable water, allowing the mixture to stand until the solid 
has settled, and then removing the pure liquid. When 
considerable calcium hydroxide is suspended in the liquid, 
the mixture is called milk of lime. Ordinary whitewash 
is thin milk of lime. Limewater is used in the chemical 
laboratory and as a medicine. 

Mortar is a thick paste formed by mixing lime, sand, and water. It 
slowly hardens or " sets/' owing to the loss of water and to the absorp- 
tion of carbon dioxide. It hardens without much shrinking, and when 
placed between bricks or stones holds them firmly in place. The sand 
makes the mass porous and thus facilitates the change of the hydroxide 
into the carbonate. The sand itself is changed chemically only to a 
slight extent, if at all. Hair is sometimes added to make the mortar 
stick better, especially when it is used as plaster for walls. 

Calcium Sulphate, CaSO 4 . Extensive deposits of the 
different forms of calcium sulphate are found in England, 
France, Nova Scotia, and in the United States, especially 
in Michigan, Kansas, Iowa, Virginia, Tennessee, and Ken- 
tucky. It is generally found in volcanic regions, and -is 
often associated with sulphur and limestone, one variety 
(anhydrite, CaSO 4 ) being found with salt. Gypsum 
occurs as white masses or transparent crystals, having the 
composition CaSO 4 . 2 H 2 O. Lustrous, translucent, soft 
crystals are called selenite. Fine grained, massive kinds 
are known as alabaster, and the fibrous kinds as satin 

Gypsum is widely used as a fertilizer, and in making 
glass and porcelain. Alabaster, being soft and beautiful, 
is carved into statues and other ornaments. 

Calcium sulphate, when heated, loses its water of crys- 
tallization, becomes opaque, and falls to a powder. This 
powder, if moistened, swells and quickly " sets " or solidi- 



Calcium, Strontium, and Barium. 327 

fies to a white, porous mass with a smooth surface. When 
properly prepared this powder is plaster of Paris, which 
derives its name from the celebrated gypsum beds near 
Paris. Plaster of Paris is used to coat walls, to cement 
glass to metal, but more largely to make casts and repro- 
ductions of statues and small objects. Stucco is essen- 
tially a mixture of glue and plaster of Paris. 

To make plaster of Paris, lumps of gypsum (CaSO 4 . 2 H 2 O) are 
heated to about 125 C. to expel part of the water. The product 
((CaSO 4 ) 2 . H 2 O) is ground fine. The " setting " is a chemical change. 
The slightly soluble plaster of Paris slowly combines with water to form 
a network of very small crystals of the less soluble hydrated calcium 
sulphate. The equation is 

(CaS0 4 ) 2 .H 2 O + 3H 2 O = 2(CaSO 4 . 2 H 2 O) 
Plaster of Paris Water Gypsum 

Calcium Compounds and Hardness of Water. Calcium 
sulphate is slightly soluble in water, and calcium carbon- 
ate, as we have already seen, is changed into the unstable 
acid carbonate by water containing carbon dioxide. Water 
containing these salts of calcium is called hard water. 
They form sticky, insoluble compounds with soap, and as 
long as water contains such salts, the soap is useless as 
a cleansing agent. Heat decomposes acid calcium carbon- 
ate, and the hardness due to calcium carbonate is called 
temporary hardness, because boiling removes it. But 
the hardness caused by calcium sulphate cannot be so re- 
moved, and is called permanent hardness. Magnesium 
sulphate, like calcium sulphate, produces permanent hard- 
ness. Soft water, such as rain water, contains little or no 

ilcium or magnesium salts. 

Calcium Chloride, CaCl 2 , is a white solid. It absorbs 
moisture rapidly, and is used to dry many gases and 

<;uids. The crystallized variety dissolves readily in 

328 Descriptive Chemistry. 

water, and the solution is attended by a marked fall of 
temperature. A mixture of crystallized calcium chloride 
and snow produces a temperature of 40 C. The liquid 
left from the interaction of calcium carbonate and hydro- 
chloric acid contains calcium chloride, which on concentra- 
tion is deposited in large crystals. These readily absorb 
water, but lose their own water of crystallization when 
heated above 200 C. This anhydrous calcium chloride is 
porous, and is the form usually used as a drying agent. 
At a high temperature it melts, and solidifies in cooling 
to a hard mass known as fused calcium chloride. 

Calcium chloride is found in small quantities in some of the Stass- 
furt salts. It is obtained in large quantities as a by-product in the 
manufacture of sodium carbonate (by the Solvay process) and other 

Other Compounds of Calcium have already been discussed and may 
be reviewed here. They are calcium fluoride, calcium carbide, the cal- 
cium phosphates, and calcium hypochlorite. Calcium sulphide (CaS) 
is formed by heating a mixture of gypsum and carbon ; like other sul- 
phides, it stains silver brown. 

Test for Calcium. Calcium compounds, especially the chloride, 
color the Bunsen flame a yellowish red. 


Strontium, Sr, and Barium, Ba, are uncommon metallic elements. 
They resemble calcium closely in their physical properties and chem: 
relations. The metals themselves never occur free, and are har !! 
more than chemical curiosities. Their compounds are abundant, and 
some are useful. 

Compounds of Strontium. The important native compounds are 
the beautifully crystallized minerals, strontianite (strontium carbor . 
SrCO 3 ) and celestite (strontium sulphate, SrSO 4 ). Strontium oxidf 
(strontia, SrO), like Ihne, is made by heating the carbonate. It unite, 
with water to form strontium hydroxide (Sr(OH) 2 ), which is used in 
the manufacture of beet sugar. Strontium nitrate (Sr(NO 3 ) 2 ) and 
other salts of strontium color a flame crimson, and are widely useo i>-. 

Calcium, Strontium, and Barium. 329 

making fireworks, especially "red fire." The latter is a mixture of 
potassium chlorate, shellac, and strontium nitrate. 

The production of the crimson colored flame is the test for stron- 

Compounds of Barium. The most abundant native compounds are 
witherite (barium carbonate, BaCO 3 ) and barite (barium sulphate, 
BaSO 4 ) . The oxides, BaO and BaO 2 , have already been mentioned as 
a source of oxygen. Barium hydroxide (Ba(OH) 2 ) solution is often 
called baryta water, and it forms the insoluble barium carbonate 
(BaCO 3 ) when exposed to carbon dioxide. Barium chloride (BaCl 2 ) is 
used in the laboratory to test for sulphuric acid and soluble sulphates, 
because it readily interacts with them and forms the insoluble barium 
sulphate (BaSO 4 ). This precipitated salt is a fine, white powder, and 
being cheap and heavy it is a common adulterant of the ordinary white 
paint, Ground native barium sulphate has a similar use. Barium sul- 
phate is also used to increase the weight of paper and to give it a gloss. 
Barium salts color a flame green, and barium nitrate (Ba(NO 3 ) 2 ) is 
extensively used in making fireworks, especially "green fire." Com- 
mercial barium sulphide (BaS), as well as the sulphides of calcium 
and strontium, shine feebly in the dark, after having been exposed to a 
bright light. On account of this property they are used in making 
luminous paint. Soluble barium salts are poisonous. 

The production of the green flame is the test for barium. 


1. Name the alkaline earth metals. What is the symbol of each ? 

2. Name several compounds of calcium. What proportion of the 
earth's crust is calcium ? 

3. Describe the preparation and state the properties of calcium. 

4. What is the formula of calcium carbonate ? State the properties, 
occurrence, and uses of (a) limestone, and () marble. 

5. State the essential characteristics of (a) calcite, () Iceland spar, 
(c) stalactites, (d) Mexican onyx, (e) travertine, (/) coquina, (g) 
chalk, (Ji) coral. 

6. Review the properties of calcium carbonate, especially its solu- 
bility (see Chapter XIV). 

7. State the uses of (a) limestone, () marble, (c) chalk. 

8. Describe the formation of (a) limestone caves, () chalk, (<:) 

jjo Descriptive Chemistry. 

9. What is the formula and chemical name of lime ? State the 
properties and uses of lime. How is it made ? State the equation for 
the chemical change. 

10. What is (a) quicklime, (fr) slaked lime, (c) hydraulic lime, (d) 
Portland cement, (e) " air-slaked " lime ? 

1 1 . What is the formula of calcium hydroxide ? How is it formed ? 
What are its properties ? How does it interact with carbon dioxide ? 
State the equation for the reaction. 

12. What is () limewater, (b) milk of lime, (c~) whitewash ? 

13. What is mortar ? How is it prepared ? For what is it used ? 
How does it change chemically with age ? What is plaster ? 

14. Discuss the occurrence of calcium sulphate. State the chief 
properties of (a} gypsam, (b} selenite, (c) alabaster, (d) satin spar. 
For what are gypsum and alabaster used ? 

15. What is plaster of Paris ? Why so called ? How is it pre- 
pared ? What is its chief property ? What are its uses ? What is 
the chemical explanation of " setting '' ? What is stugco ? 

1 6. What is hard water ? How does it act with soap ? What is 
(#) temporary hardness, and (b) permanent hardness ? How may each 
be removed ? What is soft water ? Why is rain water often called soft 
water ? 

17. Summarize the properties of calcium chloride. What is its 
formula ? How is it prepared ? 

1 8. Review the essential properties of (a) calcium fluoride, (b) cal- 
cium carbide, (c) tricalcium phosphate, (d} bleaching powder. 

19. What is the test for (a) calcium, (b) strontium, (c) barium ? 

20. State the use of (a) strontium hydroxide, and {b) strontium 

21. For what are (a) barium hydroxide, (b} barium nitrate, (V) 
barium sulphide, and (//) barium chloride used ? Describe barium 


1. What is the per cent of calcium in (a) marble (CaCO <3 ), 
() gypsum (CaSO 4 . 2 H.,O), (V) fluor spar (CaF 2 ), (d} superphos- 
phate of lime (CaH 4 (PO 4 ) 2 ) ? 

2. How many tons of limestone must be heated to produce 100 
tons of quicklime ? (Assume CaCO 3 = CaO + CO 2 .) 

3. Calculate the simplest formula of a compound having the per- 
centage composition Ca = 40, C = 12, O = 48. 


THESE elements form a natural group, though the mem- 
bers are not so closely related as the alkali and alkaline 
earth groups. Zinc and cadmium are much alike, and both 
also resemble magnesium. Mercury differs somewhat 
from zinc and cadmium, but resembles copper. 


Occurrence of Magnesium. Magnesium is never 
found free. In combination it is widely distributed and 
very abundant, constituting about 2.5 per cent of the 
earth's crust. Dolomite is magnesium calcium carbonate 
(CaMg(CO 3 ) 2 ) ; it forms whole mountain ranges and vast 
deposits; beds hundreds of feet thick cover thousands of 
square miles in the upper Mississippi valley. Dolomite 
closely resembles marble and limestone. Magnesium 
carbonate is also abundant. Many of the Stassfurt 
salts contain magnesium, for example, kainite (KC1, 
MgSO 4 . 3 H 2 O), carnallite (KC1, MgCl 2 .6 H 2 O), and 
kieserite (MgSO 4 . H 2 O). It is also a component of 
serpentine, talc, soapstone, asbestos, meerschaum, and 
other silicates. The sulphate and chloride are found in 
sea water and in mineral springs. 

Through the decay of rocks, magnesium compounds find their way 
into the soil, from which they are taken up by plants. Magnesium 
phosphates are found in the bones of animals and the seeds of grains, 
and also in guano, 

Descriptive Chemistry. 


II (7= 

Preparation of Magnesium. Magnesium was formerly prepared by 
reducing the chloride with sodium. It is now economically manufac- 
tured by electrolysis. A sketch of the essential parts of the apparatus 
is shown in Figure 67. Carnallite is put into the cylindrical iron vessel, 
C, which is the cathode. This is closed by the air-tight cover through 
which pass the pipes, Z>, D ', for conveying inert gases into and out of 
the apparatus. The carbon anode, A, dips into the carnallite and 
is inclosed by the porcelain cylinder, B, which is provided with a 

pipe, E, for the escape of the chlorine 
liberated at the anode. The carnallite 
is kept fused by external heat. When 
the current passes, the chlorine liberated 
at the anode escapes through E, and 
the magnesium liberated at the cathode 
floats on the fused carnallite and is pre- 
vented from oxidizing by the inert gas 
supplied through D. The porcelain 
cylinder, B, prevents the chlorine from 

escaping into the larger vessel. The 
FIG. 67. Apparatus 'for the 

manufacture of magnesium by the molten magnesium is carefully removed 
electrolysis of carnallite. at intervals. 

Properties of Magnesium. Magnesium is a lustrous, 
silvery white metal. It is a light metal, the specific grav- 
ity being only 1.75. It is tenacious and ductile, and when 
hot may be drawn into wire or pressed into ribbon, the 
latter being a common commercial form. It melts at a red 
heat and may be cast into different shapes. At a high 
temperature it volatilizes. It is easily kindled by a match 
or candle, and burns with a dazzling white light, producing 
dense white clouds of magnesium oxide (MgO). It does 
not tarnish in dry air, but in moist air it is soon covered 
with a film of oxide. It liberates hydrogen from acids. 
Heated in nitrogen, it forms magnesium nitride (Mg 3 N 2 , 
see Composition of Ammonia). 

Uses of Magnesium. Magnesium in the form of pow- 
der is used chiefly in taking flash-light photographs. 

Magnesium, Zinc, Cadmium, and Mercury. 333 

Small quantities are used in making fire-works ; and 
both the powder and wire are used in the chemical 

Magnesium Oxide, MgO, is a white, bulky powder. It 
is formed when magnesium burns in the air, but it is man- 
ufactured by gently heating magnesium carbonate, just as 
lime is made from limestone. It is often called magnesia, 
or calcined magnesia. The native oxide is the mineral 
periclase. Magnesia dissolves with difficulty in water, 
forming magnesium hydroxide (Mg(OH) 2 ). A mixture 
of magnesia and water, with or without magnesium chlo- 
ride, hardens on exposure to the air, and is often used as a 
cement or artificial stone. Native magnesium hydroxide 
is the mineral brucite. Like lime, magnesia withstands 
a high temperature, and is, therefore, used as the chief 
ingredient of a protective mixture for steam pipes and ves- 
sels which are subjected to great heat. Magnesia is used 
as a medicine for dyspepsia and an antidote for poisoning 
by mineral acids. 

Magnesium Sulphate, MgSO 4 , is a white solid. There 
are several crystallized varieties. The native salt kie- 
serite (MgSO 4 . H 2 O) when added to water changes into 
Epsom salts (MgSO 4 . ;H 2 O). This variety was first 
found in the mineral spring at Epsom, England. It is 
very soluble in water, and its solution has a bitter taste. 
It is extensively used as a medicine, in manufacturing sul- 
phates of sodium and potassium, as a fertilizer in place of 
gypsum, and as a coating for cotton cloth. 

Magnesium Chloride, MgCl 2 , is a white solid. It is a by- 
product in the preparation of potassium chloride. The crystallized 
salt (MgCl 2 . 6 H.,O) is very deliquescent. Magnesia mixture is a mix- 
ture of magnesium chloride, ammonium chloride, and ammonium hy- 
droxide ; it is used in chemical analysis. 

334 Descriptive Chemistry. 

Magnesium Carbonate, MgCO 3 , occurs native as magnesite, and 
combined with calcium carbonate as dolomite. The commercial salt 
known as magnesia alba, or simply magnesia, is a complex compound 
(Mg(OH) 2 , 4 MgCO 3 4 H 2 O) . Several of these complex basic carbon- 
ates are known. Many face powders consist chiefly of magnesia alba. 

It was during an investigation of magnesia alba that Black discov- 
ered carbon dioxide and showed the close relation between analogous 
compounds of magnesium and calcium. 

Miscellaneous. Besides the oxide and sulphate, other compounds 
are used as medicines. Fluid magnesia, prepared by dissolving mag- 
nesium carbonate in water containing carbon dioxide, is a mild laxative. 
Magnesium citrate has a similar action ; it is an effervescing mixture 
prepared from sodium bicarbonate, tartaric and citric acids, sugar, and 
magnesium sulphate. 


Occurrence of Zinc. Free zinc is never found. The 
ores of zinc are not numerous, but are widely distributed. 
The chief ores are zinc sulphide (sphalerite, zinc blende, 
ZnS), zinc carbonate (smithsonite, ZnCO 3 ), zinc silicate 
(calamine, H 2 Zn 2 SiO 5 ), and red zinc oxide (zincite, ZnO). 
Franklinite and willemite are ores of zinc containing 
manganese and iron. Gahnite has the composition 
ZnAl 2 O 4 . 

Zinc ores are found in Germany, Italy, France, Greece, Spain, Austria- 
Hungary, Belgium, England, and the United States. Missouri and 
Kansas contain large deposits of the sulphide, while the other ores 
occur chiefly in New Jersey. About 143,000 tons of zinc were pro- 
duced in the United States in 1902, and over 60 per cent came from 
Missouri-Kansas. This was the largest amount ever produced in a 
single year. 

Metallurgy of Zinc. Zinc is easily smelted. The ores 
are first roasted to change them into the oxide, thus 

ZnCO 3 = ZnO + CO 2 

Zinc Carbonate Zinc Oxide Carbon Dioxide 

Magnesium, Zinc, Cadmium, and Mercury. 335 

ZnS +30= ZnO + SO 2 
Zinc Sulphide Oxygen Zinc Oxide Sulphur Dioxide 

The oxide is then reduced by heating it with charcoal. 
This operation is conducted in earthenware tubes or fire- 
clay crucibles connected with iron receivers into which the 
zinc vapor passes ; at first it condenses as a powder known 
as zinc dust, somewhat as sulphur forms flowers of sul- 
phur ; but it finally condenses as a liquid, which is drawn 
off at intervals and cast into bars or plates. The impure 
zinc thus obtained is called spelter ; it is freed from carbon, 
lead, iron, cadmium, and arsenic by repeated distillation, 
often under reduced pressure. 

Properties of Zinc. Zinc is a bluish white, lustrous 
metal. Its physical properties vary with the temperature. 
At ordinary temperatures it is brittle, but at 100 150 C. 
it is soft and may be rolled into sheets and drawn into 
wire, while its specific gravity rises from 6.9 to 7.2. Zinc 
which has been rolled or drawn does not become brittle 
upon cooling. At 200 C. it again becomes brittle and 
can be easily pulverized. It melts at about 433 C. and 
boils at about 940 C. Heated in the air above its melting 
point, zinc burns with a bluish green flame, forming white 
zinc oxide (ZnO). Zinc does not tarnish in dry air, but 
ordinarily it becomes coated with a dark film. Commercial 
v zinc interacts with acids and usually liberates hydrogen. 
With hot solutions of sodium and potassium hydroxides, it 
forms zincates and liberates hydrogen, thus 

2KOH + Zn = H 2 + K 2 ZnO 2 . 
Potassium Hydroxide Zinc Hydrogen Potassium Zincate 

Pure zinc interacts with acids if in contact with a platinum 
wire, or if copper sulphate solution is added. Like copper, 

33 6 Descriptive Chemistry. 

zinc withdraws other metals (e.g. lead and mercury) from 
their solutions. 

The vapor density of zinc requires the molecular weight 67.6. Since 
the atomic weight is 65.4, a molecule of the vapor contains only one 

Uses of Zinc. Zinc in stick or plates is extensively 
used as the positive plate in electric batteries. Sheet zinc 
is used as a lining for tanks, and as the protective cover- 
ing which is placed behind and beneath stoves. Iron 
dipped into melted zinc becomes coated with zinc and is 
called galvanized iron ; it does not rust easily and is widely 
used for roofs, pipes, cornices, and water tanks. Telegraph 
wire is also galvanized. Zinc dust is used in the cyanide 
process of extracting gold and in many chemical experi- 
ments in the laboratory. Brass, German silver, and other 
alloys contain zinc (see Alloys of Copper). Antifriction 
metals, which are used for bearings, are alloys of zinc. 
Babbitt's metal, for example, contains 69 per cent of zinc, 
19 of tin, 4 of copper, 3 of antimony, and 5 of lead. 

Compounds of Zinc. Native zinc oxide is red, owing 
to the presence of manganese, but the pure oxide is white 
when cold and yellow when hot. It is formed when zinc 
burns, and is manufactured in this way or by heating zinc 
carbonate. It is often called "zinc white" or " Chinese 
white," and is used to make a white paint which is not dis- 
colored by the atmosphere. Native zinc sulphide is yel- 
low, brown, or black on account of impurities, but the pure 
sulphide is white. The latter is formed as a jelly like pre- 
cipitate when hydrogen sulphide is passed into an alkaline 
solution of a zinc salt ; it is decomposed by a mineral acid. 
Zinc sulphide is also used as a white pigment. Zinc 
sulphate is. formed by the interaction of zinc and dilute 
sulphuric acid. Large quantities are made by roasting 

Magnesium, Zinc, Cadmium, and Mercury. 337 

the sulphide in a limited supply of oxygen and extracting 
the sulphate with water. It is a white, crystallized solid 
(ZnSO 4 . 7 H 2 O), which effloresces in the air, and when 
heated to 100 C. loses most of its water of crystallization. 
The crystallized salt is called white vitriol. It is used in 
dyeing and calico printing, as a disinfectant, and as a medi- 
cine. It is poisonous, but can be safely used externally to 
relieve inflammation. Zinc chloride (ZnCl 2 ) is a white, 
deliquescent solid, prepared by dissolving zinc in hydro- 
chloric acid and evaporating the solution until a sample 
solidifies on cooling. It is used in surgery, and also as a 
constituent of a mixture for filling teeth ; large quantities 
are used to preserve wood, especially railroad ties, from 
decay, nearly 1500 tons being annually consumed for this 
purpose. Zinc hydroxide (Zn(OH) 2 ) is formed by the 
interaction of sodium or potassium hydroxide and the solu- 
tion of a zinc salt. An excess of the alkaline hydroxide 
changes the zinc hydroxide into a zincate. 

Tests for Zinc. The formation of the sulphide or hydroxide, as 
above described, serves as the test for zinc. A green incrustation is 
produced when zinc compounds are heated on charcoal and then mois- 
tened with a cobaltous nitrate solution. 

Cadmium, Cd, is an uncommon metal, frequently found in zinc ores. 
It occurs native as a sulphide (greenockite, CdS). It is white, lustrous, 
and rather soft. Its specific gravity is 8.6, and its melting point is 
about 320 C. Cadmium is a constituent of certain fusible alloys (see 
Bismuth). Wood's metal contains 12 per cent of cadmium. The most 
important compound is cadmium sulphide (CdS). This is a bright 
yellow solid, formed by adding hydrogen sulphide to the solution of a 
cadmium compound. It is used as an artist's color. Its formation also 
serves as the test for cadmium. 


Occurrence of Mercury. Native mercury is occasion- 
ally found in minute globules, but the most abundant ore 

jj 8 Descriptive Chemistry. 

is mercuric sulphide (cinnabar, HgS). The ore is mined 
in Spain, Austria, Russia, Italy, and Mexico ; in the United 
States large quantities are obtained in California, and 
deposits were recently opened in Texas. 

The annual production of the United States for several years has 
been about 1000 tons. 

Mercury has been known for ages as quicksilver. The Latin name, 
hydrargyrum, which gives us the symbol Hg, means literally " water 
silver," emphasizing the fact, so well known, that mercury looks like 
silver and flows like water. 

Preparation of Mercury. Mercury is readily prepared 
by roasting cinnabar in a current of air. Sulphur dioxide 
and mercury are formed, thus 

HgS + 2 Hg + S0 2 

Cinnabar Oxygen Mercury Sulphur Dioxide 

The sulphur dioxide is usually allowed to escape, but the 
mercury vapor is condensed by passing it into large cham- 
bers, or through pear-shaped retorts or pipes, called aludels 
(see Iodine). Crude mercury is freed from dirt and me- 
chanical impurities by pressing it through linen or chamois 
leather, but it must be distilled to separate it from dissolved 
metals, such as lead or zinc. It can also be purified by 
treatment with dilute nitric acid. Mercury is sent into 
commerce in strong iron flasks, holding about 75 pounds. 
Properties of Mercury. Mercury is a bright, silvery 
metal, and is the only one which is liquid at ordinary tem- 
peratures. It solidifies at about 39.5 C. It is a heavy 
metal, the specific gravity being 13.59. It is slightly vola- 
tile even at ordinary temperatures, and the vapor is poison- 
ous. Mercury does not tarnish in the air, unless sulphur 
compounds are present. At a high temperature, it com- 
bines slowly with oxygen to form the red oxide (HgO). 

Magnesium, Zinc, Cadmium, and Mercury. 339 

Hydrochloric acid and cold sulphuric acid do not affect it ; 
hot concentrated sulphuric acid oxidizes it, and nitric acid 
changes it into nitrates. 

The vapor density of mercury requires the molecular weight 198.72. 
Since the atomic weight is 200, a molecule of the vapor contains only 
one atom. 

Amalgams are alloys of mercury with other metals. 
They are easily prepared by mixing the constituents. 
Sometimes the union is violent as in the preparation of 
sodium amalgam. Amalgamated zinc is usually used in 
electric batteries to prevent unnecessary loss of the zinc. 
Tin amalgam is sometimes used to coat mirrors. Amal- 
gams of certain metals are used as a filling for teeth. Care 
should be taken, while handling mercury, not to let it come 
in contact with rings or jewelry, since gold amalgam is 
readily formed. 

Uses of Mercury. Mercury is used in making ther- 
mometers, barometers, and some kinds of air pumps. Its 
extensive use in extracting gold and silver has been men- 
tioned (see Amalgamation). Large quantities are used in 
preparing certain medicines and explosives (e.g. fulminating 
mercury, which is used in cartridges). 

Compounds of Mercury. Mercury, like copper, forms two classes of 
compounds the mercurous and the mercuric. Mercuric oxide (HgO) 
is a red powder, produced by heating mercury in air or by heating a 
mixture of mercury and mercuric nitrate. As we have already seen, 
mercuric oxide is decomposed by heat into mercury and oxygen. A 
yellow variety is produced by the interaction of sodium hydroxide and a 
mercuric salt, thus 

2NaOH + Hg(NO 3 ) 2 = HgO + 2 NaNO 3 + H 2 O 
Sodium Mercuric Mercuric Sodium 

Hydroxide Nitrate Oxide Nitrate 

Mercurous chloride (Hg 2 Cl 2 or HgCl) is a white, tasteless powder, 
insoluble in water. It is formed when a chloride and mercurous nitrate 

34-O Descriptive Chemistry. 

interact, but it is manufactured by heating a mixture of mercuric chloride 
and mercury. Under the name of calomel it is extensively used as a 
medicine. Mercuric chloride (HgCl 2 ) is a white, crystalline solid, solu- 
ble in water and in alcohol. It is prepared by heating a mixture of 
mercuric sulphate and common salt. It is a violent poison. The best 
antidote is the white of a raw egg. The albumen forms an insoluble 
mass with the poison, which may then be removed mechanically from 
the stomach. The common name of mercuric chloride is corrosive 
sublimate. It has strong antiseptic properties, and is extensively used 
in surgery to protect wounds from the harmful action of germs ; taxi- 
dermists sometimes use it to preserve -skins, and it has many serviceable 
applications as a medicine and disinfectant. It is usually used as a 
dilute solution (i part to 1000 parts of water). Native mercuric sul- 
phide or cinnabar (HgS) is a red, crystalline solid. When hydrogen 
sulphide is passed into a solution of a mercuric salt, mercuric sulphide 
is formed as a black powder; this variety, when heated, changes into 
red crystals. 

Vermilion is artificial mercuric sulphide. It is manufactured either 
(i) by grinding together mercury and sulphur, and treating this mass 
with caustic potash solution, or (2) by heating mercury and sulphur in 
iron pans and subliming the black mass. In both processes the product 
must be carefully ground, washed, and dried. Chinese vermilion is the 
best quality. Vermilion has a brilliant red color, and, although expen- 
sive, is widely used to make red paint. 

Mercurous Nitrate (HgNO 3 or Hg 2 (NO 3 ) 2 ) and mercuric nitrate 
(Hg(NO 3 ) 2 ) are prepared by treating mercury respectively with cold 
dilute nitric acid, and with hot concentrated nitric acid. They are 
white, crystalline solids. 


1. Name the chief native compounds of magnesium. What pro- 
portion of the earth's crust is magnesium ? 

2. Describe the manufacture of magnesium by the electrolysis of 

3. Summarize the properties of magnesium. State its uses. 

4. What is the formula and chemical name of magnesium ? How 
is magnesia formed ? State its properties and uses. 

5. Describe the different varieties of magnesium sulphate. State 
the uses of Epsom salts. 

Magnesium, Zinc, Cadmium, and Mercury. 341 

6. What is the formula of magnesium carbonate ? What is (#) 
magnesite, () dolomite, (c) magnesia alba? For what is the last sub- 
stance used ? 

7. Name the chief ores of zinc. Discuss their occurrence. 

8. Describe the metallurgy of zinc. What is (a) zinc dust, and 
() spelter ? How is zinc purified ? 

9. Summarize (a) the physical properties of zinc, and () the chem- 
ical properties. 

10. State the uses of zinc. 

11. Review the alloys of copper which also contain zinc. What 
alloys are largely zinc ? 

12. Describe native and pure zinc oxide. For what is the latter 
used ? 

13. Describe zinc sulphate. How is it formed and for what is it 

14. Describe zinc chloride. For what is it used? 

15. What are the tests for zinc ? 

1 6. State the properties and uses of (a) cadmium, and () cadmium 

17. What is the chief ore of mercury ? Where is it found ? 

1 8. What is the symbol of mercury? What is the literal meaning 
of the word from which it is formed ? 

19. Describe the preparation and purification of mercury. How is it 
transported ? 

20. Summarize the properties of mercury. 

21. What are amalgams ? Name three, and state the use of each. 

22. For what is mercury used ? 

23. Describe mercuric oxide. What historical interest has it ? 

24. Describe mercurous chloride. What is its commercial name? 
State its use. 

25. Describe mercuric chloride. What is its commercial name? 
How does it differ from mercurous chloride ? State its use. 

26. What is the formula and chemical name of cinnabar ? Describe 
cinnabar. What is vermilion ? How is it manufactured? State its 

27. What is (a) magnesia, () Epsom salts, (c) galvanized iron, 
(d) Chinese white, 0) white vitriol, (/) calomel, (g) corrosive subli- 
mate ? 

342 Descriptive Chemistry. 


1. How much magnesium will be formed by heating 100 gm. of 
potassium with magnesium chloride ? (Assume K 2 + MgCl 2 = 
Mg + 2 KC1.) 

2. What is the per cent of magnesium in (#) magnesite (MgCO 3 ), 
() dolomite (MgCa(CO 3 ) 2 ), (c} Epsom salts (MgSO 4 7 H,O ) ? 

3. What is the per cent of zinc in (a) zinc sulphate (ZnSO 4 ), (b} 
zinc sulphide (ZnS), (c) zinc chloride (ZnCl 2 ), (d) zinc oxide (ZnO) ? 

4. How much zinc sulphate can be prepared from 65 gm. of zinc ? 
From 130 gm.? From 720 gm.? 

5. How much mercury is formed by decomposing 400 gm. of cin- 
nabar ? (Assume HgS + O 2 = Hg + SO 2 .) 

6. What is the per cent of mercury in (a) mercuric oxide (HgO), 
(b) calomel (Hg 2 Cl 2 ), (c) corrosive sublimate (HgCl 2 ) ? 


Occurrence. Aluminium does not occur free in nature, 
but its compounds are numerous, abundant, and widely 
distributed. About 8 per cent of the earth's crust is 
aluminium; it is, therefore, the most abundant metal. 
Many common rocks and minerals are silicates of alumin- 
ium and other metals, e.g. feldspar and mica, which make 
up a large part of granite and gneiss. Clay and slate are 
mainly silicate of aluminium, formed by the decomposition 
of complex aluminium minerals. Corundum and emery 
are aluminium oxide (A1 2 O 3 ) more or less impure. Baux- 
ite is an hydroxide of aluminium (H 4 A1 2 O 5 ). Cryolite is a 
fluoride of aluminium and sodium (Na 3 AlF 6 ). 

Aluminium was first obtained as a fine powder by Wohler in 1827. 
Deville, in 1854, prepared it in compact form and laid the foundation 
of the industry which is being developed by Hall. 

Davy proposed the name alumium, i.e. alum + him, to emphasize the 
relation of the metal to the well-known substance, alum. The word 
alumium was changed first to aluminum and then to aluminium. Some 
authorities derive the word alumium from the Latin word alumen, or 
from alumina, the common name of aluminium oxide. 

Metallurgy. Aluminium is obtained from its oxide 
(A1 2 O 3 ) by electrolysis. In the Hall process, which is 
typical, an open, iron box lined with carbon is made the 
cathode (Fig. 68). The anode consists of carbon bars 
hung from a copper rod, which can be lowered as the car- 



Descriptive Chemistry. 

bon is consumed. The process is essentially as follows : 
the bottom of the box is covered with cryolite, the anodes 
are lowered, and the box is then filled with cryolite. The 
current is turned on, and in its resisted passage through 
the cryolite enough heat is generated to melt the cryolite. 
Pure, dry aluminium oxide is now added. This is decom- 
R . posed into aluminium 

and oxygen. The oxy- 
gen unites with the 
carbon of the anodes, 
forming carbon mo- 
noxide, which burns or 

escapes. The molten 

FlG. 68. Apparatus for the manufacture of aluminium falls to the 
aluminium by the electrolysis of aluminium 
oxide. C C C is the iron box which serves as 

the cathode. A, A, etc. are carbon anodes 
attached to the copper rod, R. 

bottom. The process 
is continuous, fresh 
aluminium oxide being 
added and the molten aluminium being drawn off at inter- 
vals. The cryolite is unchanged, and merely acts as a 
solvent for the aluminium oxide. 

The United States produced about 7,000,000 pounds of aluminium 
in 1902, and the output is annually increasing. This was all produced 
at Niagara Falls. In the Heroult process, which is used in Europe and 
involves essentially the same principle as Hall's process, the aluminium 
is produced as an alloy (usually of copper) . 

Aluminium was prepared until about 1885 by a complicated process, 
(i) Bauxite was changed into aluminium oxide free from iron by fusion 
with sodium carbonate and treatment with carbon dioxide. (2) The 
aluminium oxide was then changed into aluminium sodium chloride by 
fusion with sodium chloride and charcoal and subsequent treatment with 
chlorine. (3) This chloride was reduced by sodium, thus 

A1C1 3 + 3Na = Al + 3 NaCl 
Aluminium Sodium Aluminium Sodium 
Chloride Chloride 

Aluminium. 345 

The sodium for this operation was prepared by the Castner process (see 
Sodium), and the two industries were developed simultaneously. 

The extensive application of the electrolytic method has reduced the 
price of aluminium from about $12 a pound during 1862-1887 to about 
30 cents in 1902. 

Properties. Aluminium is a bluish white metal. It is 
very light compared with other common metals, since its 
specific gravity is only about 2.6 ; this value is one third 
that of iron. It is ductile and malleable, and is often 
sold in the form of wire and sheets ; it must be annealed 
frequently during the hammering or drawing. It is a 
good conductor of heat and electricity. Its tensile 
strength is about as great as that of cast iron. It melts at 
about 660 C., and may be cast and welded, but not readily 
soldered so as to produce a permanent joint. The cap of 
the Washington Monument is a casting of aluminium 
which weighs about eight and a half pounds. Pure alu- 
minium is only very slightly oxidized by air. Hydrochlo- 
ric acid changes it into aluminium chloride, thus 

2A1- + 6HC1 = 2A1C1 3 + 3H 2 

Aluminium Hydrochloric Aluminium Hydrogen 
Acid . Chloride 

Under ordinary conditions nitric and' sulphuric acids do 
not affect it. Sodium and potassium hydroxides change it 
into aluminates, thus 

6NaOH + 2A1 = 2 Na 3 AlO 3 + 3 H 2 
Sodium Hydroxide Aluminium Sodium Alumkiate Hydrogen 

The properties of aluminium are modified by the presence of impuri- 
ties. The usual impurities are iron, other metals, and silicon. Some 
of these, especially the iron and silicon, come from the raw products 
used in its manufacture. They tend to make the metal harder and more 
active chemically, but less malleable, ductile, and tenacious. If it were 
not for the presence of these impurities in clay, this substance would be 
a cheap and inexhaustible source of aluminium. 

346 Descriptive Chemistry. 

Uses. The varied properties of aluminium adapt it to 
numerous uses. It is made into the metallic parts of mili- 
tary outfits, caps for fruit jars, surgical instruments, cook- 
ing utensils, tubes, the framework and fittings of boats and 
air ships, telephone receivers, scientific apparatus, parts of 
opera glasses and telescopes, the framework of cameras, 
stock patterns for foundry work, and hardware samples. 
Its attractive appearance has led to its extensive use as an 
ornamental metal, both in interior decorative work and in 
numerous small objects, such as trays, picture frames, 
hairpins, and combs. Aluminium leaf is used for decorat- 
ing book covers and signs ; the powder is likewise used as 
a protective and attractive coating for letter boxes, steam 
pipes, lamp-posts, radiators, smokestacks, and other metal 
objects exposed to heat or the weather. During the last 
few years aluminium wire has come into use as a conductor 
of electricity. Large quantities of aluminium are used to 
reduce oxides, to make iron and steel more fluid, and to 
produce sounder castings. The applications of aluminium 
are constantly increasing. 

Alloys. The alloy of aluminium and copper aluminium bronze 
has been been described (see Alloys of Copper) . Magnalium is a recent 
alloy containing from 75 to 90 per cent of aluminium, the rest being 

Aluminium Oxide, A1 2 O 3 , is the only oxide of alumin- 
ium. It is often called alumina, as silicon dioxide is called 
silica. Its native forms, corundum and emery, are found 
in Massachusetts, New Jersey, Georgia, Pennsylvania, 
North Carolina, and Canada ; large quantities come from 
Asia Minor and the islands near Greece. Emery is ex- 
tremely hard, and is used in various forms powder, cloth, 
paper, and wheels to grind and polish hard metals, plate 



glass, etc. The crystallized varieties of aluminium oxide 
are usually known as corundum, and the transparent, 
colored kinds have long been prized as gems (see below). 

Alumina may be prepared by burning the metal or by heating its 
hydroxide. Thus prepared, it is a white powder, insoluble in water, 
but soluble in zfcids and in the caustic alkalies. It melts in the oxyhy- 
drogen flame, and in the electric furnace. Heating lessens its chemical 
activity. When alumina or any other compound of aluminium is heated, 
then cooled and moistened with cobaltous nitrate solution and heated 
again, the mass turns a beautiful blue color. This is a test for alu- 

Aluminium is both basic and acid, that is, with acids it forms salts, 
like aluminium chloride, while with bases it forms aluminates. 

Gems containing Aluminium. Corundum (A1 2 O 3 ) has long been 
found as crystals in Ceylon, Siam, Burma, and other places in the 
Orient. The color is due to traces of impurities, usually oxides of 
metals. The sapphire is blue, and the ruby is red. The Oriental 
topaz is yellow, the Oriental amethyst is purple, and the Oriental 
emerald is green. Montana furnishes many sapphires, the output in 
1901 being valued at $90,000. These gems may be artificially produced 
by dissolving alumina in a fused substance, adding an oxide to secure 
the desired color, and then allowing the alumina to crystallize. Spinels 
are complex compounds of aluminium. The typical or ruby spinel is 
magnesium aluminate (MgAl 2 O 4 ). It resembles the true ruby. Other 
spinels differ from the ruby spinel both in color and in composition. 
Turquoise is a complex aluminium phosphate containing traces of cop- 
per. It has a beautiful robinVegg-blue color, is compact, and may be 
worked into various shapes. Formerly turquoise came almost exclu- 
sively from Persia, but now New Mexico meets all demands. Nearly 
$120,000 worth of turquoise are mined annually m that state. Topaz 
is a complex aluminium silicate containing fluorine. It is usually pale 
yellow, and is found in many localities. Emerald is, next to diamond 
and ruby, the most precious gem. It is an aluminium silicate con- 
taining the rare element beryllium. The finest specimens have a deep 
emerald-green color and come from Colombia, South America. Garnet 
is a complex silicate of aluminium and another metal, especially cal- 
cium, magnesium, iron, or manganese. The kind used as a gem has a 
deep red color and is rather abundant. 

348 Descriptive Chemistry. 

Aluminium Hydroxide, A1(OH) 3 , is a white, jelly like 
solid formed by adding an hydroxide to the solution of an 
aluminium salt, thus 

AlClg + 3 NH 4 OH = A1(OH) 3 + 3 NH 4 C1 

Aluminium Ammonium Aluminium Ammonium 

Chloride Hydroxide Hydroxide 'Chloride 

It is insoluble in water. It interacts with strong acids 
and with alkalies (except ammonium hydroxide), forming 
respectively aluminium salts and aluminates. Thus 

A1(OH) 8 + 3 HC1 = A1C1 8 4- 3 H 2 O 
Aluminium Hydrochloric Aluminium Water 
Hydroxide Acid Chloride 

A1(OH) 3 + 3 NaOH = Na 3 AlO 3 + 3 H 2 O 

Sodium Sodium 

Hydroxide Aluminate 

Bauxite is a native aluminium hydroxide, though it contains iron 
and silicon. It resembles clay in texture and color. The vast deposits 
found at Baux, in southern France, furnish most of the raw material for 
the manufacture of aluminium, though about twenty thousand tons are 
annually obtained from our Southern states, chiefly from Georgia. 

Aluminium Sulphate, A1 2 (SO 4 ) 3 . 18 H 2 O, is a white, 
crystalline solid. The commercial salt has a variable com- 
position ; and, if pure, it dissolves readily and completely 
in water. It is extensively used in dyeing and paper 
making, and in preparing other aluminium compounds. 

Aluminium sulphate is prepared from pure clay, bauxite, or cryolite. 
If clay or bauxite is heated with sulphuric acid and then allowed to 
cool, the product is impure aluminium sulphate, known as " alum cake," 
or, if much iron is present, as " alumino ferric cake. 1 ' It is used to 
purify sewage and for other purposes where iron and the other impuri- 
ties do no harm. Purer aluminium sulphate is prepared by heating 

Aluminium. 349 

bauxite with soda ash, extracting the sodium aluminate formed with 
water, and precipitating the aluminium, as the hydroxide with carbon 
dioxide gas. The relatively pure hydroxide is then changed into sul- 
phate by treatment with sulphuric acid. The product, known as 
"concentrated alum, 1 ' has the composition expressed by the formula 
A1.,(SO 4 ) 3 . 20 H 2 O, though separate crystals contain only eighteen 
molecules of water of crystallization. By boiling cryolite with milk of 
lime, the sodium aluminate thereby formed may be changed into " con- 
centrated alum," as described above. About 50,000 tons of "con- 
centrated alum " are annually produced in the United States. 

Alum. When solutions of aluminium sulphate and potas- 
sium sulphate are mixed and concentrated by evaporation, 
transparent, colorless, glassy crystals are deposited. This 
solid is potassium alum, or simply alum. It has the com- 
position represented by the formula, K 2 A1 2 (SO 4 ) 4 . 24 H 2 O, 
or K 2 SO 4 , A1 2 (SO 4 ) 3 . 24 H 2 O, and is sometimes called a 
double salt. It is the type of a class of similar salts called 
alums, which can be formed by crystallization from a 
mixture of aluminium sulphate and an alkaline sulphate. 
Alums are very soluble in water, and their solutions have 
an acid reaction and a sweetish, puckery taste. They 
crystallize alike, and contain twenty-four molecules of 
water of crystallization. When heated, alums lose their 
water of crystallization and some sulphuric acid, and fall 
to a white powder or porous mass known as burnt alum. 
Potassium alum is the most common, but ammonium and 
sodium alums are manufactured and used. Sodium alum 
is an ingredient of some baking powders. Burnt alum 
finds application as a medicine. Alum has been largely 
displaced by " concentrated alum," but the real alum is 
still used in dyeing and printing cloth, in tanning and 
paper making, in purifying water and sewage, as a medi- 
cine, for hardening plaster, in making wood and cloth fire- 
proof, and in preparing other aluminium compounds. 

350 Descriptive Chemistry. 

Alum was known to the ancients, who used it in dyeing and tanning, 
and as a medicine. It was first manufactured in Europe, about the 
thirteenth century, from native alunite, which is an impure sulphate of 
aluminium, potassium, and iron. Alunite and alum slates or shales are 
now used to some extent, but most of the alum is made from bauxite. 
Not all alums contain aluminium. This metal may be replaced by iron, 
chromium, manganese, or similar metals, producing salhich have 
the same general properties as ordinary alum. 
formula of alums is M 2 (SO 4 ) 3 . X 2 SO 4 . 24 H 7 O, in 
aluminium, iron, chromium, etc., and X a metal (or group) like potas- 
sium, sodium, ammonium. Chrome alum (K.,Cr 2 (SO 4 ) 4 . 24 H.,0) 
belongs to this class. It is a purple, crystallized solid. The other alums 
have a limited, industrial application. * 

Alums and other aluminium salts are used as mordants 
in dyeing and calico printing. Some dyes must be fixed 
in the fabric by a metallic substance, otherwise the color 
would be easily removed. The cloth to be dyed or printed 
is impregnated or printed with the mordant, and then 
heated or treated with some substance to change the mor- 
dant into an insoluble compound. The mordanted cloth is 
next passed through a vat containing the solution of the 
dye, which unites chemically or mechanically (perhaps 
both) with the metallic compound, forming a colored com- 
pound. The latter is called a "lake"; it is relatively in- 
soluble, and cannot be easily washed from the cloth, i.e. 
it is a fast color. Aluminium acetate or "red liquor" and 
aluminium sulphate, besides alum, are used as mordants 
for cotton, linen, and wool. 

Cryolite is a white, glassy, crystallized solid. It often 
resembles clouded ice, and its name means "ice stone." 
Its composition corresponds to the formula Na 3 AlF 6 (or 
A1F 3 . 3 NaF). Small fragments melt easily, even in a 
candle flame, and color the Bunsen flame yellow. The 
only locality where it is found in commercial quantities is 



southern Greenland, which yields annually about 10,000 
tons. It is used not only in manufacturing aluminium, 
but as a source of alum and aluminium hydroxide, pure 
sodium carbonate and hydroxide, hydrofluoric acid, fluor- 
ides, and one kind of glass. 

Aluminium Chloride when pure is a white powder, but it is often a 
yellowish, crystalline mass (A1C1 3 . 6 H 2 O). It is prepared by heating 
powdered aluminium in chlorine, or by passing chlorine over a heated 
mixture of aluminium oxide and carbon. Exposed to the air, it absorbs 
moisture and gives off fumes of hydrochloric acid. It dissolves in 
water with evolution of heat, and if the solution is heated, hydrochloric 
acid is expelled, owing to the transformation of the chloride into the 
hydroxide, thus 

A1C1 3 + 3 H 2 = 3 HC1 + Al(OH), 
Aluminium Water Hydrochloric Aluminium Hy- 
Chloride Acid droxide 

This salt is used in organic chemistry. 

Clay is a more or less impure aluminium silicate, formed 
by the slow decomposition of rocks containing aluminium, 
especially feldspar. Pure feldspar is a silicate of alumin- 
ium and sodium or potassium. The products of its decom- 
position are chiefly an insoluble aluminium silicate and a 
soluble alkaline silicate. The latter is washed away. The 
aluminium silicate which remains is pure clay or kaolin. 
The latter is really a hydrous silicate, having the composi- 
tion corresponding to the formula Al 2 Si 3 O 7 , 2 H 2 O. The 
composition of clay varies, because it is seldom formed 
from pure feldspar. Most kaolin contains particles of mica 
and quartz. Ordinary clay contains many impurities, e.g. 
carbonates of calcium and magnesium, quartz, and iron 
compounds. Kaolin is a white, powdery mass. It becomes 
slightly plastic when wet, and can therefore be molded 
into various shapes. Ordinary clay is very plastic when 

Descriptive Chemistry. 

wet, more easily fused than kaolin, but shrinks consider- 
ably when dried and burned ; it also contains iron com- 
pounds, which color it gray, blue, yellow, brown, and red. 
All clays have a peculiar clayey odor when moist. 

Clay is the basis of pottery, of which there are three 
general kinds : porcelain or china, stoneware, and earthen- 

Porcelain is the finest kind. It is made by heating to a high tem- 
perature a mixture of kaolin, fine sand, and some fusible substance, such 
as feldspar, chalk, or gypsum. The mass when cool is hard, dense, white, 
and translucent (if thin) ; it is not easily corroded by chemicals (ex- 
cept fused alkalies). Although it is not very porous, its surface is 
glazed, partly for protection, partly for ornament. This is done by 
coating it with a mixture similar to that used for making the porcelain 
but more easily fused, and then heating again so that the glaze will 
penetrate the surface. Stoneware is similar to porcelain, but coarser, 
because the materials are less carefully selected and prepared, and are 
not heated to such a high temperature. The best grades can hardly be 
distinguished from porcelain, but usually stoneware is much heavier 
and thicker. The cheaper kinds are made into jars, jugs, and bottles, 
especially large ones used in acid manufactories. Crockery is a fine 
grade of stoneware, though the best crockery is much like porcelain. 
If less pure, plastic clay is used and heated to a moderate temperature, 
the product is known as earthenware. This is a large class and in- 
cludes majolica, tiles, terra cotta, jugs, flowerpots, clay tobacco pipes, 
drain pipe, and bricks. This ware is porous and is usually glazed by 
throwing salt into the baking oven just before the operation is over. 
The salt volatilizes arid forms a fusible sodium aluminium silicate upon 
the surface. Cheap bricks are made from very impure clay, and their 
red color is due to iron oxides formed from the iron compounds in the 
unburned clay. Buff bricks are 'made from clay containing little or no 
iron, and clay containing silica yields fire-clay bricks, stove linings, 
retorts, and crucibles. 


1. What is the symbol and atomic weight of aluminium ? 

2. Name several compounds of aluminium and discuss their occur- 
rence. What proportion of the earth's crust is aluminium ? 

Aluminium. 353 

3. State briefly the history of aluminium. 

4. Describe the metallurgy of aluminium by (#) the Hall process, 
() the Heroult process, (V) the older chemical method. 

5. Discuss the production and cost of aluminium. 

6. (#) Summarize the properties of aluminium. (<) State its uses. 
(V) Describe its alloys. 

7. What is the formula and chemical name of alumina ? Describe 
its preparation. State its properties and uses. 

8. State the properties and uses of corundum and emery. Review 
carborundum (see Chapter X). 

9. Name seven gems containing aluminium. Describe them. 

10. Describe aluminium hydroxide. How does it interact with 
acids and with alkalies ? 

1 1 . What is bauxite ? For what is it used ? 

12. Describe aluminium sulphate. State its properties and uses. 
How is it prepared ? What is " alum cake " ? u Alumino ferric cake " ? 
State their uses. 

13. What is ordinary alum ? How is it manufactured ? State the 
general properties and uses of alums. What is (a) "concentrated 
alum, 1 ' and (^) burnt alum ? 

14. Define a mordant. Describe its use. Name several mordants. 
What is (a) a " lake," (b) red liquor ? 

15. What is the general formula of an alum ? What is chrome 
alum ? 

16. Where is cryolite found ? State its properties and uses. What 
is its formula ? 

17. Describe the preparation and state the properties of aluminium 

18. What is clay ? How is it formed ? What is kaolin ? Describe 
(a) ordinary clay, and (6) kaolin. 

19. Describe the manufacture of (a) porcelain, () stoneware, and 
(V) earthenware. Give an example of each. What is meant by 
glazing ? 


What is the per cent of aluminium in (a) cryolite (AlNa 3 F 6 ), 
() turquoise (A1,P 2 O 8 . H r Al 2 O 6 . 2 H 2 O), (V) corundum (A1 2 O 3 ), (W) 
aluminium hydroxide (A1(OH) 3 ) ? 


TIN and lead are familiar metals. They have similar 
and useful properties, which give these metals and their 
compounds numerous applications. 


Occurrence of Tin. Metallic tin is rarely if ever 
found. Tin dioxide (cassiterite, tin stone, SnO 2 ) is the 
only available ore. It is not widely distributed, but large 
deposits are found in England (at Cornwall), Germany (in 
Bohemia and Saxony), Australia, Tasmania, and the East 
Indian Islands, especially Banca and Billiton. A small 
quantity is found, but not mined, in the United States. 

Tin is one of the oldest known metals. It is mentioned in the Pen- 
tateuch, and was obtained long before the Christian era by the Phoeni- 
cians from the British Isles, which were called Cassiterides (from the 
Greek word kassiteros, meaning tin). Many ancient bronzes contain 
tin. The alchemists called it Jupiter, and used the metal and its com- 

The Latin word stannum gives us the symbol Sn and the terms 
stannous and stannic. 

Metallurgy of Tin. If the tin ore contains sulphur or arsenic, these 
impurities must be removed by roasting. The tin oxide is then reduced 
by heating it with coal in a reverberatory furnace ; the simplest equation 
for this change is 

SnO 2 + C = Sn + CO 2 

Tin Dioxide Carbon Tin Carbon Dioxide 


Tin and Lead. 355 

The molten tin which collects at the bottom of the furnace is drawn off 
and cast into bars or masses, which are often called block tin. Usually 
it is purified by melting it slowly on a hearth, inclined so that the more 
easily melted tin will flow down the hearth and leave the metallic impuri- 
ties behind. This tin may be further purified by stirring the molten 
metal with a wooden pole, or by holding billets of wood beneath its sur- 
face. The impurities which are oxidized by the escaping gases collect 
as a scum on the surface and are removed. 

Properties of Tin. Tin is a white, lustrous metal, 
which does not tarnish easily in the air. It is soft and 
malleable, and can be readily cut and hammered. It is 
softer than zinc, but harder than lead. Its specific gravity 
is 7.3. Tin may be obtained in the crystalline form, and 
when a piece of such tin is bent it makes a crackling 
sound, which is caused by the friction of these crystals 
upon one another. It melts at about 232 C, and when 
heated to a higher temperature' it burns, forming white tin 
oxide (SnO 2 ). The physical properties of tin, like those 
of zinc, vary with the temperature. Concentrated hydro- 
chloric acid changes it into stannous chloride (SnCl 2 ); 
treated with hot concentrated sulphuric acid, it forms 
stannous sulphate (SnSO 4 ) and sulphur dioxide ; and com- 
mercial nitric acid oxidizes it, the white, solid product 
being known as metastannic acid. Zinc precipitates tin 
from its solutions as a grayish black, spongy mass, which 
is sometimes filled with bright scales. 

Uses of Tin. Tin is so permanent in air, weak acids 
(like vinegar and fruit acids), and alkalies that it is exten- 
sively used as a protective coating for metals. Ordinary 
tinware is sheet iron coated with tin. The tin plate 
(sheet tin, or simply "tin") is made by dipping very clean 
sheet iron into molten tin. Tacks, nails, and many small 
iron objects are similarly tinned. Copper coated with tin 

356 Descriptive Chemistry. 

is made into vessels for cooking, and brass coated with 
tin is made into pins. Large quantities of tin plate are 
used to cover roofs. Tinned iron does not rust until the 
tin is worn off and the iron exposed, and then the rusting 
proceeds rapidly. Tin is also hammered into thin sheets 
called tin foil, though much of the tin foil now used con- 
tains lead. Many useful alloys contain tin as an essential 
ingredient. During the last few years the annual con- 
sumption of tin has been about 75,000 pounds. 

Alloys of tin are described under COPPER. Those 
containing a minor percentage of tin are .bronze, gun 
metal, bell metal, speculum metal, type metal, anti-friction 
metals, and fusible alloys. Britannia metal contains 
about 90 per cent tin, 8 per cent antimony, and the rest 
mainly copper. It is a white metal, and was formerly 
made into tableware. White metal contains less tin and 
more antimony than Britannia, though the composition 
varies. It resembles Britannia. The harder varieties of 
'white metal are used as parts of machinery, and the softer 
kinds are made into ornaments and cheap jewelry. Pew- 
ter and solder contain varying proportions of tin and lead. 
Plumbers' solder, or soft solder, is about one third tin and 
two thirds lead. It is harder than either constituent, but it 
melts at a lower temperature. Tin amalgam is sometimes 
used to coat mirrors. 

Compounds of Tin. Tin forms two series of compounds, the stan- 
nous and the stannic. Stannic oxide (SnO 2 ) has already been men- 
tioned as the chief ore of tin, and as the product formed when tin is 
burned. The artificial oxide is faint yellow when hot and white when- 
cold. The native oxide is a brown or black, lustrous, and often crystal- 
lized solid. Irregular pebbles called stream tin occur in some localities 
near rivers. Stannous chloride (SnCl.,) is formed by the interaction 
of hydrochloric acid and tin. From the concentrated solution a green- 
ish salt crystallizes (SnCl 2 . i H.,O), known as tin crystals or salt of tin. 

Tin and Lead. 357 

Stannous chloride passes readily into stannic chloride (SnCl 4 ) when 
added to mercuric chloride solution. The simplest equation for this 
change is 

SnCl 2 + 2 HgCl 2 = SnCl 4 + Hg 2 Cl 2 
Stannous Mercuric Stannic Mercurous 

Chloride Chloride Chloride Chloride 

By an extension of the simplest idea of oxidation and reduction, the 
stannous chloride in the change is said to be oxidized to stannic chlo- 
ride, but it reduced the mercuric chloride to mercurous chloride. Stan- 
nous chloride is often used as a reducing agent and as a mordant in 
dyeing and calico printing. Crystallized stannic chloride (SnCl 4 . 5 H 2 O), 
known commercially as oxymuriate of tin, is also used as a mordant. Tin 
mordants produce brilliant colors. Sodium stannate (Na 2 SnO 3 . 3 H 2 O) 
is extensively used to prepare cotton cloth for printing. 


Occurrence of Lead. Metallic lead is occasionally 
found in small quantities. The most abundant ore is lead 
sulphide (galena, PbS). Other native compounds, formed 
by the alteration of galena, are the carbonate (cerussite, 
PbCO 3 ), the sulphate (anglesite, PbSO 4 ), and the phos- 
phate (pyromorphite, Pb 5 Cl(PO 4 ) 3 ). Lead compounds are 
widely distributed, but the source of commercial lead is 
the sulphide. 

Lead has been used by civilized people since the dawn of history. 
The Chinese have used it for ages to line chests in which tea is stored 
and transported. The Romans, who obtained it from Spain, called it 
plumbum nigrum, i.e. black lead. The symbol Pb Qomes from plumbum. 
The ancients also used lead compounds (especially the carbonate and 
red oxide) as paints and cosmetics. 

The annual production of lead has increased rapidly during the last 
few years, and in 1902 it was about 800,000 tons. This vast amount 
comes chiefly from the United States, Spain, Germany, Mexico, New 
South Wales, and England. The United States in 1902 produced 
about 250,000 tons of lead from ores found mainly in the Middle West 
(Illinois, Iowa, Wisconsin, and Missouri), Colorado. Idaho, and Utah. 

358 Descriptive Chemistry. 

Metallurgy of Lead. Lead is readily obtained from galena, (i) In 
the reduction process the ore is roasted in a reverberatory furnace until 
a part of the sulphide is changed into lead oxide and lead sulphate. 
The equations for these changes are 

2 PbS 4- 3 O 2 = 2 PbO + 2 SO 2 

Lead Sulphide Oxygen Lead Oxide Sulphur Dioxide 

PbS + 2O 2 PbS0 4 

Lead Sulphide Oxygen Lead Sulphate 

The air is then shut off and the mixture of the three lead compounds is 
heated to a higher temperature. By this operation the lead sulphide 
interacts with the other lead compounds, forming lead and sulphur diox- 
ide, thus 

2 PbS + PbSO 4 + 2 PbO = sPb + 3 SO 2 
Lead Sulphide Lead Sulphate Lead Oxide Lead Sulphur Dioxide 

(2) Ores poor in lead are sometimes reduced by roasting with iron, 
which combines with the sulphur, leaving the lead free, thus 

PbS + Fe = Pb + FeS 

Lead Sulphide Iron Lead Iron Sulphide 

(3) At Niagara Falls lead is obtained from galena by electrolysis. 
Crushed galena is made the cathode, dilute sulphuric acid is the electro- 
lyte, and the bottom of the reduction pan is the anode. The sulphur 
is changed into hydrogen sulphide, which escapes into a combustion 
chamber where its sulphur is recovered or converted into sulphuric acid. 
The lead remains in the pan as a spongy mass. The silver, which 
remains in the lead obtained by reduction, is extracted by the Parkes 
process (see Silver). 

Properties of Lead. Lead is a bluish metal. When 
scraped or cut, it has a brilliant luster, which soon disap- 
pears, owing to the formation of a film of oxide. This 
coating protects the lead from further change. It is a soft 
metal, and may be scratched with the finger nail. It dis- 
colors the hands, and when drawn across a rough surface 
it leaves a black mark. For this reason it is sometimes 

Tin and Lead. 359 

called black lead (see Graphite). Lead is not tough 
enough to be readily hammered into foil or drawn into fine 
wire, but it can be rolled into sheets. It is a heavy metal, 
its specific gravity being 11.35; w ith the exception of 
mercury, it is the heaviest of the familiar metals. It melts 
at 326 C, or about 100 higher than tin and 100 lower 
than zinc. Lead, when heated strongly in air, changes 
into an oxide (mainly the monoxide, PbO). Hydrochloric 
and sulphuric acids have little effect upon compact lead. 
Nitric acid changes it into lead nitrate (Pb(NO 3 ) 2 ). Acetic 
acid (or vinegar) and acids from fruits and vegetables 
change it into soluble, poisonous compounds ; hence cheap 
tin-plated vessels, which sometimes contain lead, should 
never be used in cooking. Zinc and iron precipitate lead 
from its solutions as a grayish mass, which often has a 
beautiful treelike appearance. 

Lead in Drinking Water. Lead is slowly changed into 
soluble compounds by water containing carbon dioxide, 
ammonia, nitrates, or chlorides. But water containing sul- 
phates or carbonates forms an insoluble coating on the 
lead, thus protecting it from further action. All lead salts 
are poisonous, and if taken into the system they will slowly 
accumulate and ultimately cause serious and dangerous 
illness. Water suspected of attacking lead should never 
be drunk after it has been standing very long in lead pipes, 
but should be allowed to flow until the pipe has been filled 
with fresh water. Sometimes the water cannot be drunk 
at all. The city of Lowell, Massachusetts, recently aban- 
doned one source of its water supply because of the rapid 
solvent action of the water upon lead pipes. 

Uses of Lead. Lead is extensively used as pipe, be- 
cause it can be made into indefinitely long pieces, which 

Descriptive Chemistry. 

can be easily bent, cut, and united (by solder). The pipe is 
made by forcing softened lead through a hole 
in a steel plate or by the apparatus shown 
in Figure 69. Lead pipe is not only used 
to convey water to and from parts of build- 
ings, but as a sheath for copper wires, both 
overhead and underground. As sheet lead 
it is used to cover roofs and to line sinks, 
cisterns, and the cells employed in many 
electrolytic processes. The lead chambers 
and evaporating pans used in manufacturing 
sulphuric acid are made of sheet lead. Shot 
and bullets are lead (alloyed with a little 
arsenic). Spongy lead is used in preparing 

inthelongcylin- , r, is forced tne plates of storage batteries. 

K'tough The A11 y s of Lead are important. Type 
metal contains 70 to 80 per cent lead ; the 

FIG. 69. Ap- 

Ing* lead' pipe! 
The molten lead 

the space, D, 

varied insL by other constituents are tin and antimony. The 
the steel rod, A. latter metal expands when it solidifies and 
makes the face of the type sharp and clear. 
Solder, pewter, and fusible alloys contain lead as an 
essential constituent (see Alloys of Tin). Small quantities 
are found in brass and bronze. 

Lead Oxides. There are three important oxides. Lead 
monoxide (PbO) is a yellowish powder known as massicot, 
or a buff-colored crystalline mass called litharge. It is 
formed by heating lead above its melting point in a cur- 
rent of air. It is made this way, though considerable is 
obtained as a by-product in separating silver from lead 
(see Cupellation). Large quantities are used in preparing 
some oils and varnishes, flint glass, other lead compounds, 
and as a glaze. Lead tetroxide (red lead, minium, 
Pb 3 O 4 ) is a red powder, 'varying somewhat in color and 

Tin and Lead. 361 

composition. It is prepared by heating lead (or lead mo- 
noxide) to about 350 C. It is used in making flint glass. 
Pure grades are made into artists' paint, but the cheap 
variety is used to paint structural iron work (bridges, 
gasometers, etc.), hulls of vessels, and agricultural imple- 
ments. It is used in plumbing and gas fitting to make 
joints tight. Orange mineral has the same composition 
as red lead, and although its color is lighter, its uses are the 
same. Lead dioxide (lead peroxide, PbO 2 ), is a brown 
powder formed by treating lead tetroxide with nitric acid. 
It is used in storage batteries. 

Lead Carbonate, PbCO 3 , is found native as the trans- 
parent, crystallized mineral cerussite. It is obtained as a 
white powder by adding ammonium carbonate solution to 
lead nitrate solution. Sodium and potassium carbonates, 
however, form basic lead carbonates, which have a compo- 
sition depending upon the temperature. The most im- 
portant of these basic carbonates has the composition 
corresponding to the formula 2 PbCO 3 . Pb(OH) 2 , and is 
known as white lead. It is a heavy, white powder which 
mixes well with linseed oil, and is used extensively as a 
white paint and as the basis of many colored paints. 

White lead is manufactured by several processes. The Dutch process 
is the oldest, having been used as early as 1622. It is essentially the 
same to-day, though many details have been improved. Perforated 
disks of lead are put in earthenware pots which have a separate com- 
partment at the bottom, containing a weak solution of acetic acid 
(about as strong as vinegar). These pots are arranged in tiers in 
a large brick building, and spent tan bark is placed between each 
tier. The building is now closed except openings for the entrance and 
exit of air and steam. The heat volatilizes the acetic acid which changes 
the lead into a lead acetate. The tan bark ferments and liberates car- 
bon dioxide, which changes the lead acetate into basic lead carbonate 
or white lead. The whole operation requires from sixty to one hun- 
dred days. The slowness is the chief objection to this process. In 


Descriptive Chemistry. 

the German process acetic acid vapor, steam, and carbon dioxide are 
forced into closed chambers in which sheets of lead are suspended. It 
requires about five weeks. In the French process basic lead carbonate 
is precipitated from a basic lead acetate by carbon dioxide. Milner's 
process is a modification of the French process. Both are quicker than 
the Dutch or German processes, but the product is not considered so 
good. An electrolytic process has recently been devised. The anode 
is lead, the cathode is copper, and the electrolyte is sodium nitrate 
solution. When the electric current is passed, (i) nitric acid is liber- 
ated at the anode, and changes the lead into lead nitrate, and (2) at 
the cathode sodium is formed, which decomposes the water, thereby 
forming sodium hydroxide. The lead nitrate and sodium hydroxide 
solutions interact, forming insoluble lead hydroxide and sodium nitrate, 

Pb(NO 3 ) 2 + 2NaOH = Pb(OH) 2 + 2 NaNO 3 

Lead Nitrate Sodium Hydroxide Lead Hydroxide Sodium Nitrate 

The sodium nitrate is left in the cell to be acted upon again, but the 
lead hydroxide is changed into lead carbonate by treatment with sodium 
bicarbonate. This process is rapid, and the product is claimed to be 
as good as white lead produced by other processes. White lead paint 
often turns dark in the air, owing to the formation of lead sulphide, 
which is black. Its extensive use is largely due to its great covering 
power, i.e. a very thin layer produces a perfectly white surface, and 
therefore less paint is required for a given area. It is often adulterated 
with zinc oxide and barium sulphate; those are white solids, but they 
are cheaper and have less covering power. 

Lead Sulphide, PbS. Native lead sulphide is the min- 
eral galena, the chief ore of lead. It resembles lead in 

FIG. 70. Galena crystals (cube, octahedron and cube, octahedron). 

appearance, but is harder and is usually crystallized as 
cubes, octahedrons, or their combinations (Fig. 70). It 

Tin and Lead. 363 

has perfect cubic cleavage, i.e. it breaks into cubes or frag- 
ments more or less rectangular. It is easily changed into 
lead by heating it alone or with sodium carbonate on char- 
coal. Lead sulphide, as prepared in the laboratory, is a 
black solid. 

Black lead sulphide is readily precipitated from a lead salt solution 
by hydrogen sulphide. Its formation is the test for lead. It is changed 
into lead chloride by concentrated hydrochloric acid and into lead sul- 
phate by concentrated nitric acid. 

Other Compounds of Lead, which are important, are the chloride, 
sulphate, nitrate, chromate, and acetate. Lead chloride (PbCl 2 ) is a 
white solid formed by adding hydrochloric acid or a soluble chloride to 
a cold solution of a lead salt. It dissolves in hot water. Lead sul- 
phate (PbSO 4 ) is a white solid, formed by adding sulphuric acid or a 
soluble sulphate to a solution of a lead salt. It is very slightly soluble 
in water, but soluble in concentrated sulphuric acid, hence crude sul- 
phuric acid often contains lead sulphate. Lead nitrate (Pb(NO 3 ) 2 ) is 
a white crystallized solid formed by dissolving lead (or better, lead mo- 
noxide) in nitric acid. When heated, it decomposes into lead oxide 
(PbO), nitrogen peroxide, and oxygen. Lead acetate (Pb(C 2 H 3 O 2 ) 2 ) 
is a white, crystallized solid formed by the action of acetic acid upon 
lead or lead oxide (PbO) . It is very soluble in water and is often 
called " sugar of lead. 1 ' 


1. Name the chief ore of tin. Where is it found? What is " stream 

2. Give briefly the history of tin. What is its symbol ? Why? 

3. Describe (a} the metallurgy of tin, and (6) its purification. 

4. Summarize the properties of tin. State its* uses. 

5. What is "tin 11 ? Block tin? Tinfoil? Tinware? Sheet tin? 
Tin plate ? 

6. Describe three alloys which contain large proportions of tin. 
Name several alloys containing a minor proportion of tin. 

7. Compare native and artificial tin oxide (SnO 2 ). 

8. What is the formula of (a} stannous chloride, and (b) stannic 
chloride? What is their chemical relation? State the use of each 
chloride. What other names has stannous chloride? 

364 Descriptive Chemistry. 

9. What is the most abundant ore of lead? Name other native 

10. Give a brief history of lead. What is its symbol? Why? 

11. Discuss the production of lead. 

12. Describe the metallurgy of lead by (a) the reduction process, 
() roasting with iron, (c) electrolysis of galena. 

13. Summarize the properties of lead. 

14. State the uses of lead. 

15. Discuss the relation of lead to water. 

1 6. What is (a) type metal, (6) solder, (c) fusible alloy? 

17. Give the name and formula of the oxides of lead. 

1 8. Describe the preparation, and state the properties and uses of 
(a) litharge, (#) red lead, (c) lead peroxide. 

19. What is white lead? Describe its preparation by (a) the Dutch 
method, and () electrolysis of sodium nitrate. 

20. State the properties and uses of white lead. 

21. What is the formula and chemical name of galena? Describe 
this mineral. Describe the corresponding artificial compound. What 
is the test for lead? 

22. Describe the following salts of lead : (a) chloride, (b) sulphate, 
(c) nitrate, (d) acetate. 


1. What is the per cent of lead in (a} galena (PbS), () cerussite 
(PbCO 8 ), (c) anglesite (PbSO 4 ), (d) lead acetate (Pb(C 2 H 3 O 2 ) 2 . 3 H 2 O) ? 

2. How much litharge may be made from 40.5 gm. of lead? (As- 
sume Pb + O = PbO.) 

3. What is the per cent of tin in (a) tinstone (SnO 2 ), (b) stannous 
chloride (SnCl 2 ), (c) stannic chloride (SnCl 4 )? 


THESE elements do not belong to the same group, but 
they have several common properties and form analogous 


Occurrence of Chromium. Metallic chromium is never 
found free. Its chief ore is an oxide (chromite, chrome 
iron ore, FeCr 2 O 4 ). Native lead chromate (crocoite or 
crocoisite, PbCrO 4 ) is less common. Traces of chromium 
occur in many green minerals and rocks, e.g. emerald and 
serpentine, and verde antique marble. 

Chromite is mined chiefly in Greece, New Caledonia, New South 
Wales, Turkey, and Canada. The total annual production is about 
30,000 tons. 

The word chromium comes from the Greek word chroma, meaning 
color, and emphasizes the fact that most chromium compounds have 
decided colors. 

Preparation, Properties, and Uses. Chromium was a rare metal 
until Moissan prepared it, in 1894, in the electric furnace. Now it is 
produced in quantities by heating a mixture of chromite and carbon in 
an electric furnace. The crude chromium is refined by fusing it with 
lime. Very pure chromium is also prepared by reducing chromic oxide 
with aluminium powder. 

Chromium is a lustrous gray metal. It takes a good polish, which is 
not removed by exposure to air. It is hard, but it can be filed and pol- 
ished without difficulty. Its specific gravity is about 6.9. It is not 
attracted by a magnet. It can be fused only in the electric furnace. 

Chromium is used to harden the steel, which is to be made into 
armor, projectiles, safes, and vaults, and parts of machines used to 


366 Descriptive Chemistry. 

crush gold-bearing quartz. This hardened steel is called chrome steel. 
The commercial form of chromium is an alloy of 65 to 80 per cert 
chromium, a little carbon, and the rest iron ; this alloy is called ferro- 

Compounds of Chromium are numerous, some are com- 
plex, many pass readily into one another, and a few have 
industrial applications. The most important are potassium 
chromate, potassium dichromate, chrome alum, and lead 

Potassium Chromate (K 2 CrO 4 ) and Potassium Dichro- 
mate (or Bichromate, K 2 Cr 2 O 7 ). These compounds are 
manufactured from chrome iron ore. The crushed ore is 
mixed with lime and potassium carbonate, and roasted in 
a reverberatory furnace ; air is freely admitted and the 
mass is frequently raked. By this operation the ore is 
oxidized into a mixture of calcium and potassium chro- 
mates. The mass is cooled, pulverized, and treated with 
a hot solution of potassium sulphate, which changes the 
calcium chromate into potassium chromate. The clear, 
saturated solution of potassium chromate is changed by 
sulphuric acid into potassium dichromate ; the latter is 
purified by recrystallization from water. Potassium chro- 
mate is a lemon-yellow, crystallized solid, very soluble in 
water. Acids change it into the dichromate, thus 

2 K 2 CrO 4 + H 2 SO 4 = K 2 Cr 2 O 7 + K 2 SO 4 + H 2 O 
Potassium Sulphuric Potassium Potassium Water 
Chromate Acid Dichromate Sulphate 

Potassium Dichromate is a red solid which forms large 
crystals. It is less soluble in water than potassium chro- 
mate. Alkalies change it into a chromate, thus 

K 2 Cr 2 O 7 + 2KOH = 2 K 2 CrO 4 + H 2 O 

Potassium Potassium Potassium Water 

Dichromate Hydroxide Chromate 

Chromium and Manganese. 367 

Potassium dichromate is used in dyeing, calico printing, 
and tanning, in bleaching oils, and in manufacturing other 
chromium compounds and dyestuffs. Its uses depend 
mainly upon the fact that it is an oxidizing agent. When 
hydrochloric acid is added to potassium dichromate, oxy-r 
gen from the dichromate withdraws hydrogen from the 
acid and liberates free chlorine, thus 

K 2 Cr 2 O 7 -f 14 HC1 = 2 KC1 + 2 CrCl 3 + 3 C1 2 + 7 H 2 O 

Potassium Di- Hydrochloric Potassium Chromic Chlorine Water 

chromate Acid Chloride Chloride 

If an oxidizable substance is present, such as organic mat- 
ter, alcohol, or a ferrous compound, it is quickly oxidized. 

Potassium chromate is also formed as a yellow mass by fusing on 
porcelain or platinum a mixture of a chromium compound, potassium 
carbonate, and potassium nitrate. When the mass is boiled with acetic 
acid to decompose the carbonate and expel carbon dioxide, and then 
added to a lead salt solution, yellow lead chromate is formed. This 
experiment is often used as a test for chromium. 

Chrome Alum, K 2 Cr 2 (SO 4 ) 4 . 24 H 2 O, is a purple, crys- 
tallized solid. It is analogous in composition and similar 
in properties to ordinary alum, but it contains chromium 
instead of aluminium. It can be prepared by mixing 
potassium and chromium sulphates in the proper propor- 
tion, or by passing sulphur dioxide into a solution of 
potassium dichromate containing sulphuric acid. The 
commercial substance is a by-product obtained in the 
manufacture of alizarine, a dye which yields magnificent 
colors. Chrome alum is used as a mordant in dyeing and 
calico printing, and in tanning. 

Lead Chromate, PbCrO 4 , is a bright yellow solid, formed 
by adding potassium chromate or dichromate to a solution 
of lead salt: It is known as chrome yellow and is used 
as the basis of yellow paint When boiled with sodium 

370 Descriptive Chemistry. 

called black oxide of manganese. When heated it yields 
oxygen ; and when heated with hydrochloric acid the two 
compounds interact, forming manganous chloride, chlorine, 
and water, thus 

MnO 2 + 4HC1 = MnCl 2 + Cl a + H 2 O 
Manganese Hydrochloric Manganese Chlorine Water 
Dioxide Acid Chloride 

It colors glass and borax a beautiful amethyst, and" is often 
added to common glass to neutralize the green color. 
Enormous quantities are used in the manufacture of oxy- 
gen, chlorine, glass, and manganese alloys and compounds. 

The manganese dioxide used in the manufacture of chlorine is recov- 
ered by the Weldon process. The impure manganous chloride solu- 
tion from the chlorine still is treated with calcium carbonate to neutralize 
free acid and precipitate any iron present. Lime is added to the clear 
solution of manganous chloride, and air is blown into the mixture. The 
manganous chloride is changed into manganous hydroxide (Mn(OH).,), 
which interacts with the oxygen (of the air) and lime, forming chiefly 
calcium manganite (CaMnO 3 , or CaO . MnO 2 ). After this mixture has 
settled, the calcium chloride is drawn off, and the manganese compound, 
which is called " Weldon mud," is used to generate more chlorine. 

Manganese dioxide was used by the ancients to decolorize glass, but 
its nature was misunderstood. They confused it with an iron oxide 
called magnesia stone, and the alchemists in the Middle Ages gave 
the name magnesia to this manganese dioxide. Later they called it 
magnesia nigra, or black magnesia, to distinguish it from magnesia alba, 
or white magnesia (MgO), supposing that the two were related. Man- 
ganese was isolated in 1774, and later was given the specific name 
manganesium, which was soon shortened to manganese. 

Potassium Permanganate, KMnO 4 , is a dark purple, 
glistening, crystallized solid, though the crystals sometimes 
appear black, with a greenish luster. It is very soluble in 
water, and the solution is red, purple, or black, according 
to the concentration. Potassium permanganate gives up 
its oxygen readily and is used as an oxidizing agent in the 

Chromium and Manganese. 371 

laboratory and on a large scale to purify stagnant water 
and sewage. It is such a powerful oxidizing agent that it 
cannot be filtered through paper, but only through asbestos 
or spun glass. It is also used as a disinfectant, as a medi- 
cine, in bleaching and dyeing, in coloring wood brown, and 
in purifying gases, such as hydrogen, ammonia, and carbon 

Potassium permanganate is manufactured by oxidizing a mixture of 
manganese dioxide and potassium hydroxide, and treating the resulting 
potassium manganate with sulphuric acid, carbon dioxide, or chlorine. 
The essential reactions are represented thus 

MnO, + 2KOH + O = K,MnO 4 + H 2 O 
Manganese Potassium Potassium 

Dioxide Hydroxide Manganate 

3 K 2 MnO 4 + 2 CO 2 = 2 KMnO 4 + K 2 CO 3 + MnO 2 

Potassium Permanganate 

The uses of potassium permanganate depend mainly upon its oxidiz- 
ing power. With sulphuric acid the action is represented thus 

2KMnO 4 + 3H 2 SO 4 = 50 + 2 MnSO 4 + K 2 SO 4 + 3 H 2 O 
Potassium Sulphuric Oxygen Manganese Potassium Water 
Permanganate Acid Sulphate Sulphate 

The liberated oxygen attacks at once any organic matter present, and 
the solution becomes brown or colorless, owing to the decomposition 
of the potassium permanganate into colorless compounds. 

Compounds of Manganese, like those of chromium, are numerous, 
often complex, and closely related. There are four oxides besides 
manganese dioxide. Three manganous compounds are important, the 
chloride (MnCL,), the sulphate (MnSO 4 ), and the sulphide (MnS). 
The chloride and sulphate are pink, crystallized salts, and the sulphide 
is a flesh-colored precipitate formed by adding ammonium sulphide to 
the solution of a manganous salt, thus distinguishing it from all other 
sulphides. Manganates are salts of the hypothetical manganic acid 
(H 2 MnO 4 ). They are analogous to chromates, and the manganese in 
them acts as a non-metal. Potassium manganate is obtained as a 
green mass by fusing a mixture of a manganese compound, potassium 

372 Descriptive Chemistry. 

hydroxide (or carbonate), and potassium nitrate. Its formation on a 
small scale constitutes the test for manganese. Sodium manganate 
is used in solution as a disinfectant. 


i . What is the symbol of chromium and of manganese ? Why is each 
element so named? 

2. What is the chief ore of chromium? Where is it found? What 
other minerals contain chromium? 

3. Describe the preparation of chromium. State its properties and 
uses. What is chrome steel? Ferrochrome? 

4. Describe the manufacture of (a) potassium chromate, and () po- 
tassium dichromate. State their properties and uses. What is the 
formula of each? 

5. What are the tests for chromium? 

6. Describe chrome alum. How is it made? State its uses. How 
does it differ from ordinary alum? 

7. Describe lead chromate. How is it formed ? For what is it used ? 

8. In what two ways does chromium act in its compounds? What 
is chromic oxide? For what is it used? What is chromium trioxide? 
How is it related to potassium dichromate? 

9. Name several ores of manganese. What is the chief ore ? Dis- 
cuss the production of manganese ores. 

10. Describe the preparation, and state the properties of manganese. 

11. What is spiegel iron? Ferromanganese? State their uses. 

12. Describe manganese dioxide. State its properties and uses. 
How is it recovered by the Weldon process? What is the common 
name of manganese dioxide? Why is it so called? 


1. What is the per cent of chromium in (a) lead chromate (PbCrO 4 ), 
() chrome ironstone (Cr 2 O 3 . FeO), (c) chromic oxide (Cr 2 O 3 ) ? 

2. What is the per cent of manganese in (<z) manganese dioxide 
(MnO 9 ), () manganese sulphide (MnS), (<:) manganese alum 
K 2 Mn 2 (S0 4 ) 4 .2 4 H 2 0)? 

3. How much manganese ore containing 85 per cent of manganese 
dioxide is needed to prepare 300 Ib. of chlorine? (Assume MnO 2 + 
4HC1 = C1 2 + MnCl, + 2 H 2 O.) 


Introduction. These three elements form a natural 
group. Their properties are similar. Cobalt and nickel 
are very closely related and are seldom found alone. Iron 
resembles manganese and chromium. 


Iron is the most useful of all metals. It has been known 
for ages, and has been indispensable in the development of 
the human race. 

The symbol of iron, Fe, is from the Latin wordferrum. Yromferrum 
are derived the forms ferri- and ferro- (found in such words as ferricya- 
nide, ferro manganese, ferrocyanide, etc.), and the terms ferrous and 

Occurrence of Iron. Uncombined iron is found only 
in meteorites, which fall upon the earth from remote 
regions in space, and in a very few rocks. Combined iron 
is abundant and widely distributed. It is found in most 
rocks and many minerals, in the soil, in springs and nat- 
ural waters, in chlorophyll the green coloring matter of 
plants, and in haemoglobin the red coloring matter 
of blood. The chief ores of iron are hematite (Fe 2 O 3 ), 
limonite (Fe 2 O 3 . Fe 2 (OH)) 6 , magnetite (Fe 3 O 4 ), and sider- 
ite (FeC0 3 ). 

Other abundant compounds of iron not used as a source of the metal 
are pyrites (FeS 2 ), pyrrhotite (varying from Fe 6 S 7 to Fe n S 12 ), and the 
copper-iron sulphides (chalcopyrite, CuFeS 2 , and bornite, Cu 3 FeS 3 ). 


372 Descriptive Chemistry. 

hydroxide (or carbonate), and potassium nitrate. Its formation on a 
small scale constitutes the test for manganese. Sodium manganate 
(NaMnOJ is used in solution as a disinfectant. 


1 . What is the symbol of chromium and of manganese ? Why is each 
element so named ? 

2. What is the chief ore of chromium? Where is it found? What 
other minerals contain chromium? 

3. Describe the preparation of chromium. State its properties and 
uses. What is chrome steel? Ferrochrome? 

4. Describe the manufacture of (#) potassium chromate, and () po- 
tassium dichromate. State their properties and uses. What is the 
formula of each? 

5. What are the tests for chromium? 

6. Describe chrome alum. How is it made? State its uses. How 
does it differ from ordinary alum ? 

7. Describe lead chromate. How is it formed ? For what is it used ? 

8. In what two ways does chromium act in its compounds? What 
is chromic oxide? For what is it used? What is chromium trioxide? 
How is it related to potassium dichromate? 

9. Name several ores of manganese. What is the chief ore ? Dis- 
cuss the production of manganese ores. 

10. Describe the preparation, and state the properties of manganese. 

11. What is spiegel iron? Ferromanganese ? State their uses. 

12. Describe manganese dioxide. State its properties and uses. 
How is it recovered by the Weldon process? What is the common 
name of manganese dioxide? Why is it so called? 


1. What is the per cent of chromium in (a} lead chromate (PbCrO 4 ), 
(b) chrome ironstone (Cr 2 O 3 . FeO), (<r) chromic oxide (Cr 2 O 3 ) ? 

2. What is the per cent of manganese in (a} manganese dioxide 
(MnO 2 ), (<) manganese sulphide (MnS), (c) manganese alum 
K 2 Mn 2 (SO 4 ) 4 .24H,0)? 

3. How much manganese ore containing 85 per cent of manganese 
dioxide is needed to prepare 300 Ib. of chlorine? (Assume MnO 2 + 
4HC1 = C1 2 + MnCl, + 2 H 2 O.) 


Introduction. These three elements form a natural 
group. Their properties are similar. Cobalt and nickel 
are very closely related and are seldom found alone. Iron 
resembles manganese and chromium. 


Iron is the most useful of all metals. It has been known 
for ages, and has been indispensable in the development of 
the human race. 

The symbol of iron, Fe, is from the Latin word fer 'rum. Yromferrum 
are derived the forms ferri- and ferro- (found in such words as ferricya- 
nide, ferro manganese, ferrocyanide, etc.), and the terms ferrous and 

Occurrence of Iron. Uncombined iron is found only 
in meteorites, which fall upon the earth from remote 
regions in space, and in a very few rocks. Combined iron 
is abundant and widely distributed. It is found in most 
rocks and many minerals, in the soil, in springs and nat- 
ural waters, in chlorophyll the green coloring matter of 
plants, and in haemoglobin the red coloring matter 
of blood. The chief ores of iron are hematite (Fe 2 O 3 ), 
limonite (Fe 2 O 3 . Fe 2 (OH)) 6 , magnetite (Fe 3 O 4 ), and sider- 
ite (FeC0 3 ). 

Other abundant compounds of iron not used as a source of the metal 
are pyrites (FeS 2 ), pyrrhotite (varying from Fe fi S 7 to Fe n S 12 ). and the 
copper-iron sulphides (chalcopyrite, CuFeS 2 , and bornite, Cu 3 FeS 3 ). 



Descriptive Chemistry. 

The United States leads the world in the production of iron ore, the 
annual output for the last few years being over 25,000,000 tons. This 
vast quantity comes from twenty-five different states, but the bulk is 
mined in Minnesota, Michigan, Alabama, Wisconsin, Tennessee, Vir- 
ginia and West Virginia, and Colorado. The most abundant ore is the 
red hematite, which comes chiefly from the Lake Superior region (Fig. 
71) ; large quantities are mined in Alabama and Tennessee. The 
latter states, together with Virginia and West Virginia, furnish most of 
the limonite or brown iron ore. Pennsylvania, New Jersey, and New 
York contribute most of the magnetite, though some is mined also in 

FIG. 71. Deposits of iron and copper near Lake Superior. No. 4 is the cop- 
per region. The iron regions, known as ranges, are Marquette (i), Menominee 
(2), Gogebic (3), Vermilion (5), Mesabi (6). 

Michigan. The carbonate ores, which constitute less than one per cent 
of the output, come mainly from Ohio, Maryland, and New York. Im- 
provements in the machinery and methods used in mining and trans- 
porting iron ore have reduced its cost and facilitated its production. 
Thus, at an incredibly small expense, ore from the Lake Superior region 
is raised from open pits by steam shovels, dumped into large cars, car- 
ried to shipping ports on the lakes, dumped again into huge bunkers, 
dropped down chutes into big freight steamers (many of which hold 
6000 tons), which carry it to South Chicago and Milwaukee, though 
over two thirds is received at ports on the south shore of Lake Erie 

Iron, Nickel, and Cobalt. 


and forwarded by rail to Pittsburg, Pennsylvania. This city is the great 
center of the iron and steel industries. Birmingham, Alabama, is the 
center of the industry in the South, because near it the necessary ore, 
coal, and limestone are con- 
veniently located. 

Metallurgy of Iron. 

Iron is extracted most 
easily from its oxides. 
The ores, whatever their 
character, are first 
crushed and roasted to 
change them into ferric 
oxide (Fe 2 O 3 ) as far as 
possible, and to make 
the raw material porous. 
Thus prepared, the ore 
is smelted with coke (or 
coal) and limestone in 
a blast furnace. The 
carbon reduces the oxide 
to metallic iron, which 
collects as a liquid at the 
bottom of the furnace 
beneath the slag formed 
by the limestone and 
impurities. The blast 
furnace (Fig. 72) is a 
huge tower, from forty 
to ninety feet high and 
from fourteen to seven- 
teen feet in diameter at 
the largest part; but it 
is narrower at the top 

FIG. 72. Blast furnace. A, throat; B, 
bosh ; C, crucible where the melted iron col- 
lects ; D, pipes for hot air blast ; E, escape 
pipe for gases which do not escape through 
the " down comer " ; G, cup ; //, cone ; N, 
trough for drawing off slag ; T, tuyere ; /, hole 
through which iron is withdrawn. 

and bottom than in the middle. 

376 Descriptive Chemistry. 

It is built of masonry and iron, and lined with fire brick. 
Pipes at the bottom, called tuyeres, allow large quantities 
of hot air to be forced into the furnace and up through the 
contents, thereby producing the high temperature required 
in the melting ; while another pipe at the top not only per- 
mits the escape of hot gaseous products, but conducts 
them into a series of pipes which lead to different parts of 
the plant, where the hot gases are utilized as fuel. The 
blast pipes correspond to the bellows used by a blacksmith, 
and the exit pipe to a chimney, except that gases escaping 
through chimneys are usually wasted. 

When the furnace has been heated to the proper tem- 
perature, or is already in operation, the ore, coke, lime- 
stone, etc., are carried to the top of the furnace by 
machinery and introduced into the furnace by dumping 
them upon the cone-shaped cover; their weight lowers 
the cover, which flies back tightly into place after the 
materials roll into the furnace. The charge consists of 
alternate layers of ore, fuel, and flux. The fuel is coke, or 
coke mixed with coal. The flux varies with the ore, but 
it is usually limestone, though feldspar and sand are used 
if the ore contains lime compounds. The object of the 
flux is twofold, (i) It removes the impurities from the 
charge in the form of a fusible glass called slag or cinder, 
and (2) thereby prevents the reduced iron from reuniting 
with oxygen of the air which is being constantly blown in. 
As the smelting proceeds, the iron falls through the slag 
to the bottom of the furnace, where both are drawn off 
through separate openings. Fresh charges of definite 
weight are added at regular intervals, and the whole oper- 
ation continues without interruption for months or even 

The iron from the furnace is usually poured into molds 

Iron, Nickel, and Cobalt. 377 

of sand and allowed to solidify. Such iron is called pig 
iron or cast iron. A single large furnace will produce in 
a day about 500 tons of pig iron. In some plants the 
molten iron is run into huge vessels, called converters, and 
made directly into steel (see below). 

The chemical changes involved in the metallurgy of iron are numer- 
ous and complicated. In general, the iron oxide is reduced to metallic 
iron largely by carbon monoxide. The carbon of the fuel at first forms 
carbon dioxide with the oxygen of the air blast. But the dioxide is 
soon reduced by the hot carbon to the monoxide, which interacts with 
the ore, thus 

Fe 2 s + 3 CO = 2Fe + 3 CO 2 
Ferric Carbon Iron Carbon 

Oxide Monoxide Dioxide 

Considerable carbon monoxide escapes, however, in the waste gas. At 
this stage the iron becomes porous, and is doubtless prevented from re- 
oxidation by the carbon dioxide liberated from the decomposed lime- 
stone. As the spongy iron sinks into the hotter part of the furnace, it 
combines with carbon to some extent, finally melts, sinks through the 
slag, and accumulates at the bottom, or hearth, of the furnace. The 
iron obtained in this way contains small amounts of carbon, sulphur, 
phosphorus, silicon, and manganese. 

About 300 blast furnaces were in operation in the United States in 
1902, and the number is increasing. They consumed over 30,000,000 
tons of ore, 18,000,000 tons of fuel (chiefly coke), and 9,500,000 tons of 
limestone. They produced nearly 18,000,000 tons of pig iron over 
one third of the world's output. Germany and the United Kingdom 
produced the bulk of the remainder. 

Varieties of Iron. The iron we use and speak of is 
not pure iron, but largely a mixture or compound of iron 
with other elements, chiefly carbon. It is customary to 
speak of three varieties of iron, cast iron, steel, and 
wrought iron. This classification is based chemically 
upon the per cent of carbon they contain, though their 
physical properties are modified by the presence of silicon, 

37 8 Descriptive Chemistry. 

phosphorus, sulphur, and manganese. Each typical vari- 
ety has specific properties ; but the different varieties are 
closely related, and pass easily and gradually into each 
other. Commercially, there are several kinds of cast iron 
and many kinds of steel. 

Cast Iron is the most impure variety. It contains, be- 
sides carbon, the impurities mentioned above. It has a crys- 
talline structure, and is brittle. The proportion of carbon 
varies from 1.5 to 6 or more per cent. If most of the car- 
bon is combined with the iron, the metal is called white 
cast iron. But if the molten metal cools slowly, much of 
the carbon remains uncombined as graphite, and the color 
of the iron is gray; this kind is gray cast iron. It is 
softer than the white variety, and melts at a lower tempera- 
ture. Although cast iron is brittle, it will withstand great 
pressure. It cannot be welded or forged, that is, hot 
pieces cannot be united, nor be shaped by hammering. 
But it is extensively used to make castings. This is the 
kind of iron used in an ordinary iron foundry. The iron, 
which melts at a comparatively low temperature (about 
1100 C.), is heated in a furnace similar to a blast furnace, 
and when molten is poured into sand molds of the 
desired shape. Stoves, pipes, pillars, railings, parts of 
machines, and many other useful objects are made of cast 
iron. Birmingham, Alabama, is the center of the cast-iron 
industry in the United States. 

Cast iron containing much manganese is called spiegel iron or ferro- 
manganese (see Manganese). About 300,000 tons of this kind are 
annually produced in the United States. 

Wrought Iron is the purest variety of commercial iron. 
It contains not more than 0.5 per cent of carbon and some- 
times only 0.06 per cent, the average being o. 1 5 per cent. 
It is tough, malleable, and fibrous. It can be. bent. Un- 

Iron, Nickel, and Cobalt. 379 

like cast iron, it does not withstand pressure, but it will 
sustain great weight. An iron wire will sustain the weight 
of nearly a mile of itself. It melts at such a high tem- 
perature (1600 to 2000 C.) that it is not used for casting. 
It can be forged and welded, and is therefore often called 
malleable iron. It may be seen undergoing these opera- 
tions in a blacksmith's shop. It can also be rolled into 
plates and sheets and drawn into fine wire ; in these forms 
the metal is very strong. Wrought iron is made into wire, 
sheets, rods, nails, spikes, bolts, chains, anchors, horseshoes, 
tires, and agricultural implements. It is less important 
than formerly, since it is being largely replaced by steel. 

Wrought iron is made from cast iron by burning out the 
impurities. The process is technically called puddling. 
Cast iron is heated in a furnace, much like a reverberatory 
furnace, lined on the bottom and sides with iron ore (fer- 
ric oxide, Fe 2 O 3 ). The intense heat melts the cast iron ; 
its carbon and silicon are removed partly by the oxygen of 
the air, but mainly by the oxygen of the iron oxide. As 
the mass becomes pasty, owing to its higher melting point, 
it is stirred vigorously, or " puddled." At the proper time 
the lumps are removed and hammered, or more often 
rolled between ponderous rollers. This operation removes 
the slag, and if the rolling is repeated, the quality of the 
iron is improved ; the final rolling leaves the iron in the 
shape desired for market. 

Steel is intermediate between cast iron and wrought iron 
as far as its proportion of carbon is concerned. Many 
grades of steel are manufactured, and their physical prop- 
erties depend not only upon the presence of other elements 
besides carbon, especially phosphorus, silicon, and certain 
metals, but also upon the raw materials, the method of 
manufacture, and subsequent treatment. 

380 Descriptive Chemistry. 

The properties of steel are numerous. It is both 
fusible and malleable, and hence can be forged, welded, 
and cast. It is harder, stronger, and more durable than 
pure iron, and is more serviceable. But its most valuable 
property is the varying hardness which it can be made to 
acquire. If steel is heated very hot and then suddenly 
cooled by immersion in cold water or oil, it becomes brittle 
and very^hard. But if heated and then cooled slowly, it 
becomes soft, tough, and elastic. All grades of hardness 
may be obtained between these extremes. And if the 
hardened steel is reheated to a definite temperature, deter- 
mined by the color the metal assumes, and then properly 
cooled, a definite degree of hardness and elasticity is ob- 
tained. This last operation is called tempering. Every 
kind of tool has a temper determined by its use. Special 
grades of hard steel are also made by the addition of cer- 
tain metals, especially chromium and nickel. Harveyized 
steel is made by packing steel in a mixture of charcoal 
and boneblack, and heating it to a very high temperature. 
This operation hardens the surface. This brand of steel 
is extensively used as armor plate in warships. 

Manufacture of Steel. The aim in the manufacture 
of steel is to prepare a product containing little or no sul- 
phur, phosphorus, and silicon, but the desired proportion 
of carbon. This may be done by three general methods : 

(1) the carbon, may be partly removed from cast iron, 

(2) carbon may be added to wrought iron, (3) cast iron 
may be added to wrought iron. The first method is diffi- 
cult to operate, and is seldom used. The other methods 
are utilized by several processes. 

(i) In the cementation or crucible process, wrought 
iron and carbon are packed in tight fire-clay boxes and 
heated for several days, The iron slowly absorbs carbon 

Iron, Nickel, and Cobalt. 


in some way unknown at present, and becomes a steel of 
extreme purity and excellent quality. The bars are melted 
in graphite crucibles to make the metal of uniform quality, 
and cast into large bars called ingots. This process is 
long and expensive, but the steel is considered the best for 
fine tools. 

(2) The Bessemer process is the one in most general 
use. It was devised in about 1860, and has practically 
revolutionized steel making. By the economical, scien- 
tific, and extensive application of this process, all grades of 
steel are quickly made at such a relatively small cost that 
the use of this metal has been enormously extended, much 
to the prosperity of the United States. About two thirds 
of the annual production is Bessemer steel. The process 
consists in burning out the impurities in cast iron by forc- 
ing air through the molten metal, and then adding just 
enough cast iron (spiegel iron) to give the desired propor- 
tion of carbon. The operation is 
carried on in a converter (Fig. 73). 
This is a huge, egg-shaped vessel, 
supported so that it can be rotated 
into different positions; it is also 
provided with holes at the bottom 
through which a powerful blast of 
air can be blown. It is made of thick 
wrought iron plates, and is lined with 
an infusible mixture rich in silica. 
The converter is swung into a hori- 

FlG. 73. Converter. 

zontal position and five to twenty tons of molten pig iron 
are poured in direct from the blast furnace. The air blast 
is turned on, and the converter is swung back to a vertical 
position. As the air is forced through the molten metal, 
the temperature rises, the carbon is oxidized to carbon 

382 Descriptive Chemistry. 

monoxide which burns on the surface of the metal, and the 
silicon is oxidized to silicon dioxide, which is taken up by 
the slag. This oxidation generates enough heat to keep 
the metal melted, and no fuel need be used. As soon 
as the impurities have been burned out, sufficient spiegel 
iron is added to change the wrought iron into steel. By 
adding spiegel iron of known composition, Bessemer steel 
of any desired grade is produced. After the completion 
of the operation, which takes about twenty minutes, the 
contents of the converter are poured into molds. 

(3) In the Bessemer process, sulphur and phosphorus 
are not removed. Both are objectionable impurities ; sul- 
phur makes steel brittle when hot, and phosphorus makes 
it brittle when cold. The Thomas-Gilchrist process is a 
modification of the Bessemer process by which the sulphur 
and phosphorus can be removed. The converter is lined 
with a mixture of lime and magnesia, called a basic lining, 
lime is also added to the charge of pig iron, and the blast 
is continued a little longer than in the Bessemer process, 
otherwise the operations are the same. The phosphorus 
forms a phosphate and the sulphur a sulphate, both of 
which are taken up by the lining. The lining, which is 
known as Thomas slag, is used as a source of phosphorus 
for fertilizers. 

(4) In the Siemens-Martin or open-hearth process a 
mixture of cast iron and wrought iron (or steel) in proper 
proportions is melted on a hearth with an oxidizing gas 
flame. Old wrought iron or cast iron, known as "scrap," 
can be used. When a test shows that the metal contains 
the desired proportion of carbon, ferromanganese is added, 
and the charge is then poured into molds. This process 
requires a special furnace and gas plant, and is more ex- 
pensive than the Bessemer process, since it takes longer. 

Iron, Nickel, and Cobalt. 383 

But it is easily controlled, and yields a tough, elastic steel, 
which is excellent for bridges, large machines, large guns, 
and gun carriages. Immense, quantities of the nickel steel 
used for the armor plate are made by the open-hearth 
process. The production of the open-hearth steel has more 
than doubled in the last few years, being over five and 
a half million tons in 1902. 

Uses of Steel. Steel is now used instead of iron for 
many purposes. High buildings, bridges, rails, cars, loco- 
motives, battleships, electrical machinery, boilers, agricul- 
tural implements, wire nails, rods, hoops, tin plates, and 
castings of all kinds consume vast amounts. Its extensive 
use in making springs, tools, cutlery, pens, needles, etc., 
need not be further mentioned. 

Properties of Iron. Chemically pure iron, though un- 
known in commerce, may be obtained in the laboratory 
by reducing the oxide or chloride with hydrogen or with 
alcohol. Such iron is called iron " by hydrogen," or " by 
alcohol." The purest commercial form is the wrought 
iron used for piano wire. Pure iron is a silvery white, 
lustrous, metal. It is softer than ordinary iron, but melts 
at a higher temperature. The specific gravity is about 
7.8. It is attracted by a magnet, but soon loses its own 
magnetism. Dry air has no effect upon iron, but moist 
air containing carbon dioxide rusts it. Iron rust is a com- 
plex compound, but its essential constituent is a ferric 
hydroxide (Fe 2 O 3 . Fe 2 (OH) 6 ). Rusting proceeds rapidly, 
because the film of rust is not compact enough to protect 
the metal. Like many metals, iron readily interacts with 
dilute acids, and as a rule hydrogen and v ferrous com- 
pounds are the products. 

384 Descriptive Chemistry. 

With nitric acid various products result, according to the conditions, 
ferrous nitrate and ammonium nitrate, if the acid is cold, but ferric 
nitrate and oxides of nitrogen if the acid is warm. If a clean iron wire is 
dipped into fuming nitric acid and then into ordinary nitric acid, no action 
is apparent. The iron is said to be passive. This peculiar fact has not 
been adequately explained. Steam and hot iron interact, thus 

3Fe + 4H 2 O = Fe 3 O 4 + 4H 2 
Iron Water Iron Oxide Hydrogen 

(See Preparation of Hydrogen.) 

Compounds of Iron. Iron forms two series of com- 
pounds, the ferrous and the ferric. They are analogous 
to cuprous and cupric, mercurous and mercuric com- 
pounds. Ferrous compounds in an acid solution pass into 
the corresponding ferric compound by the action of oxi- 
dizing agents, e.g. oxygen, nitric acid, potassium chlorate, 
potassium permanganate, and chlorine. Conversely, ferric 
compounds are reduced to the ferrous by reducing agents, 
e.g. hydrogen, hydrogen sulphide, sulphur dioxide, and 
stannous chloride. The passage from one series to the 
other occurs easily, especially from ferrous to ferric. In 
most of its compounds, iron acts as a metal. Many com- 
pounds of iron have industrial importance, as well as 
scientific interest. 

Oxides and Hydroxides of Iron. Iron forms three 
oxides. Ferrous oxide (FeO) is an unstable black powder. 
Ferric oxide (Fe 2 O 3 ) occurs native in many varieties as 
hematite the most abundant ore of iron. It may be pre- 
pared by heating ferrous sulphate or ferric hydroxide. 
Large quantities are obtained as a by-product in the manu- 
facture of Nordhausen (or fuming) sulphuric acid and of 
galvanized iron and tinned ware. It is sold under the 
names rouge, crocus, and Venetian red. It is used to pol- 
ish glass and jewelry, and to make red paint. Ferrous- 

Iron, Nickel, and Cobalt. 385 

ferric or ferroso-ferric oxide (magnetic oxide of iron, 
Fe 3 O 4 ) occurs native as magnetite ; if noticeably magnetic, 
it is called loadstone. It is produced as a black film or 
scale by heating iron in the air ; heaps of it are often seen 
beside the anvil in a blacksmith's shop. The firm coating 
of this oxide formed by exposing -iron to steam protects 
the metal from further oxidation. 

Ferrous hydroxide (Fe(OH) 2 ) is a white solid formed by the inter- 
action of a ferrous salt and an alkali, such as sodium hydroxide. Ex- 
posed to the air, it soon turns green, and finally brown, owing to the 
formation of ferric hydroxide. Ferric hydroxide (Fe 2 (OH) 6 ) is a red- 
dish brown solid, formed by the interaction of ammonium hydroxide 
(or any alkali) and a ferric salt. Several ferric hydroxides are known. 
The freshly prepared compound is an antidote for arsenic. 

Ferrous Sulphate (FeSO 4 ) is a green salt obtained by 
the interaction of iron (or ferrous sulphide) and dilute sul- 
phuric acid, and is a by-product in several industries (e.g. 
see Ferric Oxide). It is prepared on a large scale by oxi- 
dizing iron pyrites (FeS 2 ); this is accomplished simply by 
roasting, or more often by exposing heaps of pyrites to 
moist air. The mass is extracted with water containing 
scrap iron and a small proportion of sulphuric acid. From 
the clear solution, large light green crystals are obtained. 
The crystallized salt (FeSO 4 . 7 H 2 O) is also called green 
vitriol or copperas. Exposed to the air, ferrous sulphate 
effloresces and oxidizes. Large quantities are used as a 
mordant and a disinfectant, and in manufacturing ink, 
bluing, and pigments. Much black writing ink is made 
essentially by mixing ferrous sulphate, nutgalls, gum, and 
water. Blue ink is usually made of Prussian blue an 
iron compound (see below) oxalic acid, and water. 

Ferric Sulphate (Fe 2 (SO 4 ) 3 ) is formed by oxidizing an acid solution 
of ferrous sulphate with nitric acid. When ferric sulphate solution is 

j 86 Descriptive Chemistry. 

mixed with the proper quantity of potassium (or ammonium) sulphate, 
iron alum (K 2 Fe 2 (SO 4 ) 4 . 24 H 2 O) is formed. It is a violet, crystallized 
solid, which has properties like ordinary alum. Iron alum is used 
chiefly as a mordant. 

Iron Sulphides. There are two iron sulphides. Com- 
mercial ferrous sulphide (FeS) is a black, brittle, me- 
tallic-looking solid, but the pure compound is yellow and 
crystalline. It is also obtained as a black powder by the 
interaction of a dissolved ferric or ferrous salt and ammo- 
nium (or potassium) sulphide. It is made on a large scale 
by fusing a mixture of iron and sulphur. It is used chiefly 
in preparing hydrogen sulphide. Ferric sulphide (iron 
disulphide, iron pyrites, pyrite, FeS 2 ) is one of the com- 
monest minerals. It is a lustrous, metallic, brass-yellow 
solid. Crystals of pyrites, found in many rocks, are often 
mistaken for gold hence the popular name "fool's gold." 
It is valueless as an iron ore, but large quantities are used 
as a source of sulphur in making sulphuric acid. Over one 
and a half million tons are annually consumed in the acid 
industry. The largest pyrite producers are Spain, France, 
Portugal, Germany, and the United States. The domestic 
output comes chiefly from Virginia, Colorado, Massachu- 
setts, and New York. 

Iron Chlorides. When iron interacts with hydrochloric acid, fer- 
rous chloride (FeCl 2 ) is formed in solution. Heated in the air, or 
better with potassium chlorate or nitric acid, it is changed into ferric 
chloride, thus 

2FeCl 2 + 2HC1 + O = 2FeCl 3 + H 2 O 
Ferrous Chlo- Hydrochloric Oxygen Ferric Chlo- Water 
ride Acid ride 

Ferric chloride is a black, lustrous, crystalline solid ; but owing to its 
extreme deliquescence, it is usually sold as a solution, which is a dark 
brown liquid. It is prepared by passing chlorine into a ferrous chloride 

Iron, Nickel, and Cobalt. 387 

solution, or by the interaction of iron and aqua regia. When treated 
with nascent hydrogen or another reducing agent, ferric chloride is 
changed into ferrous chloride. 

Ferrous Carbonate (FeCO 3 ) occurs native as the iron ore siderite, 
clay iron stone, or spathic iron ore. The typical variety is light yellow 
or brown, lustrous, crystalline, and not very hard ; but many kinds are 
impure, and the properties vary. It is slightly soluble in water contain- 
ing carbon dioxide, and is therefore found in some mineral springs (see 
Chalybeate Waters). Like all carbonates, it yields carbon dioxide 
with warm hydrochloric acid. 

Iron Cyanides. Iron and cyanogen (CN), with or with- 
out potassium, form several compounds. The most impor- 
tant is potassium ferrocyanide (K 4 Fe(CN) 6 ). It is a 
lemon-yellow, crystallized solid, containing three molecules 
of water of crystallization. Unlike most cyanogen com- 
pounds, it is not poisonous. Its commercial name is 
yellow prussiate of potash. It is manufactured by fusing 
together iron filings, potassium carbonate, and nitrogenous 
animal matter (such as horn, hair, blood, feathers, and 
leather). The mass is extracted with water, and the salt 
is separated by crystallization. In Germany this salt is 
manufactured from the iron oxide which has been used 
to purify illuminating gas. Large quantities are used in 
dyeing and calico printing, and in making bluing and 
potassium cyanogen compounds. Potassium ferricyanide 
(K 3 Fe(CN) 6 ) is a dark red, crystalline solid, containing no 
water of crystallization. It is often called red prussiate 
of potash. It is manufactured by oxidizing potassium 
ferrocyanide with chlorine, thus 

K 4 Fe(CN) 6 + Cl = K 3 Fe(CN) 6 + KC1 
Potassium Ferro- Chlorine Potassium Ferri- Potassium 
cyanide cyanide Chloride 

It is very soluble in water, forming a deep yellow, unstable 
solution. In alkaline solution it is a vigorous oxidizing 

388 Descriptive Chemistry. 

agent, and therefore finds extensive use in dyeing. It is 
also used as one of the ingredients of the sensitive coating 
of " blue print " paper. 

Ferrous salts and potassium ferricyanide interact in solution and pro- 
duce ferrous ferricyanide (Fe 3 (Fe(CN) 6 ) 2 ) . This is a blue solid and is 
often called TurnbulPs blue. But ferrous salts produce with potassium 
ferrocyanide a white precipitate (ferrous ferrocyanide), which quickly oxi- 
dizes to a complex blue compound. Ferric salts interact with potassium 
ferrocyanide and produce ferric ferrocyanide (Fe 4 (Fe(CN) 6 ) 2 ). This 
is a dark blue solid, and is called Prussian blue or Berlin blue. Ferric 
salts produce no precipitate with potassium ferricyanide. Prussian blue 
is extensively used in dyeing and calico printing, and in making bluing. 
The above reactions, which allow ferrous and ferric salts to be distin- 
guished, may be summarized as follows : 





Whitish precipitate 
Turnbuirs blue 

Prussian blue 
No precipitate 

Besides the above tests, potassium sulphocyanate produces a dark red 
liquid with ferric salts, but leaves ferrous salts unchanged. The tests 
for iron are thus numerous and specific. 


Nickel, Ni, occurs combined with arsenic, sulphur, or 
both. Small amounts of metallic nickel are found in me- 
teorites. The chief ores are nickel-bearing iron sulphides, 
which are abundant in the Sudbury district, Canada, and 
the silicates found in New Caledonia. A small amount is 
produced in the United States as a by-product in smelt- 
ing lead ores from a Missouri mine. 

Nickel is obtained from its Ores by complicated pro- 
cesses, and is now refined by electrolysis. It is a white 

Iron, Nickel, and Cobalt. 389 

metal, which takes a brilliant polish. It is ductile, hard, 
tenacious, and does not tarnish in the air. Like cobalt, it 
is attracted by a magnet. 

Nickel has varied Uses. For many years it has been 
used as one ingredient of the small coins of several coun- 
tries. The per cent of nickel varies from 12 in the United 
States cent to 25 in the five-cent piece. German silver 
contains from 15 to 25 per cent of nickel, the rest being 
copper and zinc. Large quantities of nickel are used to 
coat or plate other metals, especially iron and brass. The 
nickel plating is done by electrolysis, as in the case of 
silver and gold plating, though the electrolytic solution 
used is a sulphate of nickel and ammonium, not a cyanide. 
The deposit of nickel is hard, brilliant, and durable. 
Nickel becomes malleable, if a little magnesium is added 
to the molten metal, and sheets of iron covered with such 
nickel are made into vessels for cooking. Nickeloid is a 
nickel-plated sheet zinc. Its attractive appearance and 
non-corrosive property adapt it for the manufacture of 
reflectors, refrigerator linings, bath tubs, show cases, and 
signs. The most important use of nickel is in the manu- 
facture of nickel steel. This contains about 3.5 per cent 
of nickel. Large quantities are used for the armor plates 
and turrets of battleships,, and for parts of machinery 
requiring great strength. 

Nickel forms two series of compounds, the nickelous and the nick- 
elic. The nickelous are more common, and many of them are green. 
The test for nickel is the formation of the apple-green hydroxide 
(Ni(OH) 2 ) by the interaction of an alkali and the solution of a nickel 

Cobalt, Co, generally occurs combined with arsenic and sulphur, and 
is often associated with nickel compounds. It is a lustrous metal with 
a reddish tinge, harder than iron, but less magnetic. The hydrated 

390 Descriptive Chemistry. 

compounds are red in solution, anhydrous compounds are blue. 
Hence red crystallized salts turn blue when heated. Some cobalt 
compounds are used to color glass, porcelain, and paper, especially a 
cobalt silicate. This is known as smalt, or smalt blue ; and since it is 
unchanged by sunlight, acids, or alkalies, it is used to decorate porce- 
lain. Other pigments are cobalt blue (an oxide of cobalt and aluminium) , 
and Rinmann's green (an oxide of cobalt and zinc). The blue color 
produced by fusing cobalt compounds into a borax bead is the test for 


1 . What is the symbol of iron ? From what word is it derived ? 

2. Discuss the occurrence of iron. Name the chief ores. Name 
other compounds of iron. What proportion of the earth's crust is iron? 

3. Discuss () the production and transportation of iron ore in 
the United States, and ($) the production of iron. 

4. What is the general chemical change in the metallurgy of 
iron? Describe a blast furnace. Summarize the smelting of iron. 
Discuss the chief physical and chemical changes involved in the smelting. 

5. Name the varieties of iron. How do they differ essentially? 
What is (a) galvanized iron, (b) meteoric iron? 

6. Describe cast iron. State its composition, properties, and uses. 

7. Describe the manufacture of wrought iron. State its composi- 
tion, properties, and uses. 

8. State the composition and properties of steel. Compare briefly 
with cast and wrought iron. What is tempering? 

9. Describe the manufacture of steel by the following processes : 
(a) cementation, (b} Bessemer, (c) Thomas-Gilchrist, (d} Siemens- 

10. State the uses of steel. 

1 1 . State the properties of iron. 

12. How are ferrous changed into ferric compounds, and vice versa ? 

13. How is ferric oxide prepared? What is the native form called? 
For what is crocus used? 

14. What is the formula and chemical name of magnetic oxide of 
iron? What is loadstone? How is magnetic oxide of iron produced? 
What is the native form called? 

15. Describe ferric hydroxide. What is its use? 

16. Describe ferrous sulphate. How is it prepared? For what is 
it used ? What is copperas ? 

Iron, Nickel, and Cobalt. 391 

17. What is iron alum ? How is it related to ordinary alum? 

1 8. Describe ferrous sulphide. How is it made? For what is it 
used? Compare it with ferric sulphide. Discuss the occurrence and 
use of the latter. 

19. Describe ferrous carbonate. 

20. Describe potassium ferrocyanide. How is it made? State its 
properties and uses. What is its common name? Its formula? 

21. Describe potassium ferricyanide. For what is it used? How 
is it related chemically to potassium ferrocyanide? 

22. Describe the tests for iron. What is Prussian blue ? For what 
is it used? 

23. Discuss the occurrence of nickel. State its properties and uses. 
Describe nickel plating. What is (a) a "nickel," (b) nickel steel? 
What is the test for nickel ? 

24. State the properties of cobalt. For what are its compounds 
used? What is smalt? What is the test for cobalt ? 


1. Calculate the percentage composition of () ferric oxide, (b) fer- 
rous sulphate, (V) ferrous sulphide (FeS). 

2. If 1.586 gm. of iron form 2.265 K m - f ferric oxide, what is 
the atomic weight of iron? (Equation is 2 Fe + 3 O = Fe 2 O 3 .) 


Occurrence of Platinum. Platinum occurs as the essen- 
tial ingredient of platinum ore or so-called native platinum. 
The ore contains from 60 to 86 per cent of platinum. The 
other metals present are ruthenium, osmium, iridium, rho- 
dium, and palladium. Iron, gold, and copper are also usu- 
ally present. Only one native compound is known, viz. 
platinum arsenide (sperrylite, PtAs 2 ). 

The ore is found chiefly in the Ural Mountains in Russia, but some 
comes from South America, Australia, and Borneo. The United States 
produced about 1400 ounces of metallic platinum in 1901 the largest 
annual output on record. It came from the gold deposits in California 
and the copper mines in Wyoming. The latter source also furnished 
osmium, palladium, and iridium. The world's annual production of 
metallic platinum for the last few years has been about 165,000 ounces. 
Russia supplies over 90 per cent of this amount. 

The word platinum is derived from platina, a form of the Spanish 
word plat a, meaning silver, because native platinum was regarded as an 
impure ore of silver by the Spaniards, who first discovered it in South 
America about 1735. Platinum is now sometimes called by its old 
name platina. 

Preparation of Platinum. The platinum ore, which occurs in 
rounded grains or flattened scales, is first digested with dilute aqua 
regia to remove the gold, silver, and copper ; and then with concen- 
trated aqua regia, which changes all the platinum and a very little 
iridium into soluble compounds, leaving behind an alloy of iridium and 
osmium. From the clear solution the platinum and iridium are precipi- 
tated by ammonium chloride as compounds, which, on heating, yield 
the metals as a spongy mass. This spongy platinum is melted in a 


Platinum and Associated Metals. 393 

lime crucible with an oxhydrogen flame, or hammered while hot into 
sheet platinum. The very small amount of iridium is seldom removed 
from the metallic platinum. 

Properties and Uses of Platinum. Platinum is a lus- 
trous, grayish white metal. It is malleable and ductile, 
and usually appears in commerce in the form of wire and 
sheets. Sheet platinum is cut into squares the familiar 
platinum foil of the laboratory, or made into crucibles, 
dishes, Vid stills for sulphuric and hydrofluoric acid (Fig. 
74). Its use in these forms is due partly to its infusibility 
and partly to its resistance to acids and other corrosive 
chemicals. Although it is attacked by fused caustic alka- 
lies and a few 
other substances, 
it is practically 
indispensable in 
the chemical lab- 
oratory. Plati- 
num is a good 

FIG. 74. A platinum dish. 

conductor or elec- 
tricity, and large quantities are consumed in incandescent 
electric light bulbs. Short pieces of wire are fused 
into the glass at the base of the bulb and attached to. 
the outside wires, conveying the current to and from the 
carbon filament within. Platinum is the only metal thus 
far found which is perfectly adapted to this use. Dentists 
use alloys of platinum as a filling for teeth, and some is 
made into jewelry. The demand exceeds the supply, and 
in the last five years the price of this rare metal has 
doubled, being $21 an ounce in 1902. Platinum has a 
specific gravity of about 21, which is higher than that of 
any known substance, except osmium and iridium. In the 
form of a black, porous mass it is called spongy platinum, 

394 Descriptive Chemistry. 

and a still finer form is called platinum black. Both forms 
absorb large volumes of gases ; and if a current of the gas 
is directed against the metal, the gas often takes fire. Me- 
tallic platinum has the same property to a less degree, for 
it becomes red-hot if held in a stream of illuminating gas, 
and often ignites the gas. Palladium has similar proper- 
ties (see Occlusion). Platinum forms alloys with other 
metals, and should never be heated with lead, similar met- 
als, or their compounds, since the alloys have a low melt- 
ing point. With iridium, however, it forms a very hard 
alloy of which the international metric apparatus is made. 

Platinic Chloride (PtCl 4 ) is the only important compound of plati- 
num. It is a brownish solid formed by treating platinum with aqua 
regia and evaporating the solution to dry ness. The solution is used in 
chemical analysis, and in photography to produce " platinum prints. 1 ' 
Chloroplatinic acid (H 2 PtCl 6 ) forms complex salts, of which the yel- 
low, crystalline potassium chlorplatinate (K 2 PtCl 6 ) and ammonium 
chlorplatinate ((NH 4 ) 2 PtCl 6 ) are the best known. 

The Metals associated with Platinum have limited uses. Pal- 
ladium is used in chemical analysis to absorb hydrogen, osmium is 
utilized in the Auer incandescent electric light, and a native (as well as 
an artificial) alloy of iridium and osmium, called iridosmine, is used to 
tip gold pens. 


i. Name the metals related to platinum. 
2.. Discuss the occurrence of platinum. 

3. What is (a) native platinum, (b) spongy platinum, (c) platinum 
black, (d) platinum foil, (e) sheet platinum ? 

4. Discuss the production of platinum. 

5. What is the symbol of platinum ? What is the derivation of the 
word platinum ? 

6. Describe the preparation of platinum. Summarize its properties. 
State its uses. 

7. Describe platinic chloride. 

8. State the uses of the metals related to platinum. 

Platinum and Associated Metals. 395 


1. A piece of platinum foil measuring 10.5 cm. by 1.5 cm. weighs 
0.723 gm. Into how many pieces, each weighing i dg., may it be 
divided ? 

2. The specific heat of platinum is 0.0324. According to analysis, 
35.5 gm. of chlorine unite with 48.6 gm. of platinum to form platinic 
chloride. What is (a) the atomic weight of platinum, and (b) the 
formula of platinic chloride ? 


Introduction. In the preceding chapters emphasis has 
been laid on individual elements. Certain group relations 
were also pointed out, but little or nothing was said con- 
cerning the elements as a single large group. The ele- 
ments are not independent. They possess certain funda- 
mental properties, which show that although apparently 
very different, they are really closely related. In this 
chapter we shall consider two topics which illustrate this 
general fundamental relationship, viz. the periodic law and 
spectrum analysis. 


Classification of the Elements. As the number of 
elements increased, attempts were made to classify them. 
About the time of Lavoisier (1743-1794) they were roughly 
divided into metals and non-metals. Those elements 
were called metals which were hard, lustrous, heavy, and 
good conductors of heat, while the others were called non- 
metals. This classification proved to be misleading as 
additional elements were discovered. It is used, however, 
even now, because many common elements fall readily into 
one of these classes. 

Classification according to acid and basic properties 
prevailed for a time. But it was abandoned largely be- 
cause such a basis of division excluded elements exhibiting 


General Relations of the Elements. 397 

both acid and basic properties, such as arsenic, antimony, 
chromium, and aluminium. 

The elements have also been classified according to their 
valence into six or seven groups (the mono-, di-, tri-, etc.). 
But this plan has been largely given up on account of so 
many troublesome cases of variable and unsatisfied valence 
(see Valence). 

About 1828 Dumas pointed out striking resemblances 
between certain elements, and he suggested several groups 
or families. For example: 









This classification was arbitrarily based on selected physi- 
cal and chemical properties. It was interesting but incom- 
plete, because it emphasized resemblances and overlooked 
differences that is, the basis of comparison was not 
broad enough. 

The first actual progress began to be made about 1850, 
when chemists became deeply interested in the significance 
of atomic weights. Dumas (in 1857) and others pointed 
out certain remarkable numerical relations existing be- 
tween the atomic weights of related elements. Thus, the 
atomic weight of sodium is half the sijm of the atomic 
weights of lithium and potassium 

Li = 7, Na = 2 3> K = 39. ^~ = 2 3 - 
The same is true of phosphorus, arsenic, and antimony 
P= 3 i, As= 75> Sb=i20. 3I + I2 = 7S . 5 . 

398 Descriptive Chemistry. 

The existence of other relations similar to these, together 
with a deep desire to obtain more accurate atomic weights 
and a growing interest in the properties of the elements 
themselves, focused the attention of chemists at this time 
(1855-1865) upon the relation of properties to atomic 
weights. Several things fostered the above principle. 
One was the atomic weight determinations of Stas, whose 
masterly work proved beyond doubt that Prout was incor- 
rect when he insisted in 1815 that the atomic weights are 
whole numbers. Another was the acceptance by most 
chemists of the same table of atomic weights. A third 
was the rapid accumulation of many facts about the ele- 
ments and their compounds. Chemists were ready for a 
new classification of the elements. 

The Periodic Classification. Previous to 1869 no 
classification included all the elements. In that year the 
Russian chemist Mendeleeff published a classification of 
the elements according to the periodic law. His views 
had been partially anticipated by several chemists, and 
were soon amplified by the German chemist, Lothar Meyer. 
Their classification of the elements revealed a new relation 
between the properties of the elements and their atomic 
weights. If all the elements are arranged in the order of 
their increasing atomic weights beginning with lithium, 
their properties will vary periodically, i.e. at certain regu- 
lar intervals or periods elements will be found which have 
similar properties. In other words, a certain increase in 
atomic weight causes a reappearance or return of prop- 
erties. The general relation is often summarized in the 
Periodic Law - 

The properties of the elements are periodic functions of 
their atomic weights. 

General Relations of the Elements. 




, in ' 

^ " co vo 1 w ro ^ 

iT IT !T if j i | ill 




? CO g 2 

II - II 8 II II 1 1 
"1 1 ~ 1 1 1 


^ ! ^ 

2 1 S 1 1 ~ 1 I 




. . o? 

4- H in N 
M ro f* M w> S 


^ ^ IT < r ^ ii ; s 

> ' 6 H 1 


10 ON vd 

5- f 1 | 1 l ' ' . * ? 

H N U 1 H 


C4 t^. M O* 

M f ^ f -5 f ' | P 

' co > i3 W 1 



v? S ' 

a .* . *;-. 1.- S 4-..I i i 

, 1, | N W 

o s d ^ p I i 




hx ts. 

^COOs^^ilcol (if 

iTnir 3 7 W) ii' 3 

a i M ^ s * ^ L_i_J 


H M ^1 Or^QO Oi OH 

400 Descriptive Chemistry. 

Function here means the exhibition of some special rela- 
tion, viz. that of properties to atomic weight. Interpreted 
freely, the law means (i) properties and atomic weight are 
related, they depend upon each other; and (2) this relation 
is exhibited again and again as we reach elements with 
increasing atomic weights at regular intervals in the suc- 
cessive arrangement 

The Periodic Table originally proposed by Mendeleeff 
has been modified from time to time, as new facts have 
necessitated. The table generally accepted at the present 
time is given on page 399. 

From the table it is seen that the elements fall naturally 
into two subdivisions, (i) Those in the same vertical col- 
umn belong to the same natural group or family. Thus, 
in Group I are found the alkali metals, in Group II the 
alkaline earth metals, in Group VII the halogens. (2) 
The elements in the same horizontal row belong to the 
same period. The periodic variation of their properties is 
well illustrated by the second and third periods. Begin- 
ning with lithium, the general chemical properties vary 
regularly with increasing atomic weight Thus, the metal- 
lic character gradually diminishes until fluorine is passed 
and sodium is reached; here it reappears. Proceeding 
onward from sodium, the same gradation of properties is 
noticed until potassium is reached, and here again the 
marked metallic character in the same way reappears. 
There is no sudden change in properties until we pass 
from one period to the next. Thus, fluorine at the end of 
the second period forms a powerful acid, but sodium at the 
beginning of the third period forms a strong base. Simi- 
larly, chlorine is strongly acidic ; but potassium, which 
begins the next period, is markedly basic; chlorine is a 
typical non-metal, while potassium is a typical metal. Not 

General Relations of the Elements. 401 

all elements fit the periodic classification equally well, but 
the arrangement is at least very suggestive, and doubtless 
expresses an approximately truthful relation. 

The Gaps in the Periodic Classification probably corre- 
spond to elements not yet discovered. Three such gaps, 
which were in the original table, have been filled. When 
Mendeleeff proposed his arrangement, he predicted the 
discovery of three elements having definite properties. 
These elements, gallium, scandium, and germanium, 
have since been discovered and now occupy their pre- 
dicted place in the table. Possibly other gaps will be filled 
by newly discovered elements. 

The discovery of the predicted elements was not the only immediate 
service of MendeleefFs table. It also emphasized the necessity of more 
accurate atomic weights. Several elements did not fall into their proper 
places, and careful investigation showed that their accepted atomic 
weights were incorrect. Thus, the atomic weights of beryllium and in- 
dium were changed to their present values, and the present order of the 
platinum metals was adopted ; cobalt and nickel are still being studied. 
The position of argon, helium, and very rare metals is still doubtful, 
owing to a limited knowledge of their properties and atomic weights. 
Hydrogen, also, still lacks a place. 


Introduction. When light from an ordinary gas flame, 
glowing lime or other solid, or a Welsbach flame is passed 
through a prism and falls upon a white surface, a long 
band of color is produced. The colors are perfectly 
blended, and are arranged like the familiar colors of the 
rainbow. This band of colors is called a spectrum. The 
white light has been separated or analyzed into the col- 
ored. The examination and study of the spectrum of a 
substance is spectrum analysis, and it is accomplished by 
a spectroscope. 

402 Descriptive Chemistry. 

The Spectroscope consists essentially of a prism and tubes, one of 
which is a telescope (Fig. 75) . The light enters a slit in the tube, passes 

FIG. 75. A spectroscope. 

through, and falls upon the prism. Here it is bent from its path, and as 
it emerges from the prism, it may be viewed through the telescope as a 
magnified spectrum. 

Kinds of Spectra. (i) The spectrum of an incan- 
descent solid is a continuous band of colors. (2) But the 
spectra of gases are narrow, colored, vertical bars or lines, 
separated by black spaces. Thus, sodium vapor has a yel- 
low line, potassium a red and a violet line, and barium sev- 
eral lines where the green and yellow parts of the ordinary 
spectrum occur. Each element which is a gas, or which can 
be vaporized, has its own bright line spectrum. The lines 
always occupy the same relative positions, which in most 
cases have been very carefully determined. Therefore, 
when examined through a spectroscope, the yellow line of 
sodium will always be seen in its proper place, and the red 
and violet potassium lines in their places. Therefore, by 
examining the light from different substances, it is possi- 

General Relations of the Elements. 403 

ble to tell what elements they contain. (3) The spectrum 
of sunlight is the familiar band of colors, but it is crossed 
vertically by many black lines, which have fixed positions 
(Fig. 76). It is believed that the sun is a glowing hot 
solid, surrounded by very hot gases. It therefore should 

AaBC D Eb F G If 

111 ill I 

Hed Orange Fellow Green Blue Indigo Violet 

FIG. 76. Spectrum of sunlight showing some of the vertical lines. 

give the two kinds of spectra, the continuous and the 
bright line. Now it has been proved that the vapor of 
an element absorbs the light given out by the same ele- 
ment when solid. Hence the dark lines which appear in 
the solar spectrum are caused by the absorptive power of 
the gases in the sun's atmosphere. The solar spectrum is 
often called an absorptive spectrum. 

Spectrum Analysis. In the laboratory the spectro- 
scope is used to detect the presence of certain elements, 
more especially the metals. If the metal or one of its 
compounds is put on a platinum wire and held in the 
Bunsen flame before the slit, the characteristic spectrum 
of the element can be easily recognized in the telescope. 
Two spectra do not interfere, because each line has its 
own place. Hence several elements may be distinguished 
in a mixture. Minute quantities are easily detected by the 
spectroscope. Rare elements, which can be obtained only 
in very small quantities or with great difficulty, are studied 
by the spectroscope. Thus, Bunsen, who (with Kirch- 
hoff) devised the improved spectroscope, discovered the 
rare metals, rubidium and caesium. And within the last 
few years the spectroscope has been especially serviceable 

404 Descriptive Chemistry. 

in studying argon, helium, krypton, neon, and xenon. 
By means of the spectroscope it has been shown that the 
sun contains many elements found in our earth. Accord- 
ing to a reliable authority, about thirty of the elements 
known to us are present in the sun. The spectroscope 
also enables astronomers to tell the nature of stars, comets, 
nebulae, and other heavenly bodies. The stars thus far 
examined give spectra crossed by dark lines, and therefore 
these bodies are like the sun ; but nebulae give bright line 
spectra, and hence consist of incandescent gases. 


1. Discuss the classification of the elements according to (a) metals 
and non-metals, () acid and basic properties, (c) valence, (d) groups 
based on resemblances, (e) numerical relations. 

2. What is the fundamental idea of the periodic classification ? 
How does it differ from previous systems ? When and by whom was 
this classification proposed and developed ? 

3. State the periodic law. Explain it. What is meant by (a) func- 
tion, () period, (c} group ? 

4. Illustrate the law by (a) the alkali metals, and (b} the halogens. 

5. Discuss the gaps in the periodic arrangement of the elements. 

6. Of what use has this law been ? 

7. State some objections to it. 

8. Describe (a) a continuous spectrum, () a line spectrum, (c) an 
absorption spectrum. 

9. Describe a spectroscope. How is it used ? 

10. What kind of a spectrum is produced by (a} a glowing solid, 
() a glowing vapor, (c) a glowing solid surrounded by a glowing vapor? 

n. What is spectrum analysis ? How is it applied (a) in the labo- 
ratory, and (b) by astronomers ? 

12. What does spectrum analysis show about each element ? About 
their relations to each other ? About their distribution ? About the 
heavenly bodies ? 

13. Who perfected the spectroscope and developed its use ? 

14. What recent use has been made of the spectroscope in (a) chem- 
istry, and (b) astronomy ? 


Introduction. In the early days of chemistry it was 
believed that starch, sugar, and other compounds obtained 
from plants and animals were produced by the influence of 
some mysterious vital force. Such compounds were called 
organic, because of their connection with living things, i.e. 
with bodies having organs ; and they were sharply dis- 
tinguished from inorganic or mineral compounds obtained 
from the earth's crust. This distinction prevailed until 
Wohler, in 1828, prepared urea a characteristic organic 
compound from inorganic substances. Since then the 
barrier between the two classes of compounds has been 
completely removed. We now believe that compounds of 
carbon, whatever their source, are subject to the laws that 
govern all other compounds. The terms organic and inor- 
ganic are still used, though they have lost their original 
narrow meaning. Carbon forms a vast number of com- 
pounds which are related to each other, and which differ 
markedly from most compounds of other elements. It is 
convenient, therefore, to distinguish these compounds by 
the term organic and to study them under the comprehen- 
sive title of Organic Chemistry or the Chemistry of Carbon 

Composition of Organic Compounds. The number of 
organic compounds is very large, but they contain only a 
few elements seldom more than four or five. Hydro- 


406 Descriptive Chemistry. 

carbons, as already indicated, contain carbon and hydro- 
gen. Vegetable substances, typified by starch, sugar, and 
fruit acids, contain carbon, hydrogen, and oxygen. Ani- 
mal substances, like hah", albumen, gelatine, and muscle 
generally contain nitrogen as well as carbon, hydrogen, 
and oxygen ; some also contain sulphur or phosphorus. 
Artificial organic compounds, like dyestuffs, may contain 
any element, especially chlorine, iodine, and metals. 

The number and complexity of organic compounds is 
due to several facts already mentioned in a previous 
chapter, (i) Atoms of carbon have power to unite with 
themselves. (2) Atoms of different elements can be intro- 
duced into carbon compounds. Sometimes these atoms 
are simply added, sometimes they replace other atoms, 
thus producing an endless number of addition and substi- 
tution products. (3) The same number of atoms may 
arrange themselves differently, thereby producing isomeric 
compounds having different properties. To these princi- 
ples, which should be reviewed until firmly grasped, must 
be added another. (4) Organic compounds contain radi- 
cals. These radicals are analogous to hydroxyl (OH) and 
ammonium (NH 4 ), and like these radicals they exist only 
in combination. They act like single atoms and enter 
unchanged into a number of organic compounds. The 
radical C 2 H 5 is called ethyl. It is present in many 
organic compounds, and its presence in ordinary alcohol 
gives rise to the scientific name, ethyl alcohol. Methyl 
(CH 3 ) is another important radical, and phenyl (C 6 H 5 ) is 
especially common in the benzene series of organic com- 

Structure of Organic Compounds. An extensive study 
of the properties of organic compounds has revealed many 
facts about their constitution, i.e. the structure of their 

Some Common Organic Compounds. 407 

molecules. Little or nothing, of course, is known about 
the shape, size, etc., of molecules, but much is known 
about the grouping of atoms and of radicals in the mole- 
cules. These facts, which are ascertained by experiment 
and are often too complex to be expressed briefly, may be 
represented by suitable formulas. The ordinary or empiri- 
cal formula of alcohol is C 2 H 6 O. But this formula tells 
nothing about the relation these atoms bear to each other, 
nor whether all the hydrogen atoms act alike. Experiment 
proves, however, that (i) one hydrogen atom acts differ- 
ently from the other five, and (2) one hydrogen atom is 
always associated with the oxygen atom in chemical 
changes. Hence, the formula C 2 H 5 . OH expresses more 
fully these facts. Such a formula is called a rational or 
constitutional formula. Sometimes constitution is ex- 
pressed by a graphic formula. Thus methane and ethane 
have the graphic formulas 

H H H 

I I I 

H C H H C C H 

I I I 

H H H 

Methane Ethane 

In these diagrams the single lines represent a valence of 
one nothing else, and the number of lines connected 
with each atom must be equal to the valence of the ele- 
ment in the compound. The lines are sometimes called 
bonds or links, but they are not intended to represent at- 
traction or any other force. Nor do they represent space 
relations. In the case of methane, they mean that the 
four hydrogen atoms bear the same relation to the single 
carbon atom. In the case of ethane, they mean the same, 

408 Descriptive Chemistry. 

but they also indicate that the two carbon atoms are joined. 
The graphic formula of ethyl alcohol is 

H H 

I I 
H C C O H 

I I 
H H 

This is not an arbitrary arrangement ; the facts mentioned 
above necessitate this general arrangement. Additional 
illustrations of this subject will be given, as different 
compounds are discussed. 

Classification of Organic Compounds. Organic com- 
pounds are divided and subdivided into many classes 
for purposes of study. Only the most common organic 
compounds can be considered in this book. These are 
members of the following groups: (i) Hydrocarbons, (2) 
Alcohols, (3) Aldehydes, (4) Ethers, (5) Acids, (6) Ethe- 
real salts, (7) Fats, glycerine, and soap, (8) Carbohydrates, 
(9) Benzene and its derivatives. Some compounds are so 
closely related that they really belong to several of these 
groups, while a few cannot strictly be put in any of them. 


Three of these compounds of carbon arid hydrogen have 
been fully considered in Chapter XV. 7 The chief facts 
and fundamental principles recorded there may be profit- 
ably reviewed at this point. Other hydrocarbons will be 
discussed under Benzene (see below). 


Alcohols are compounds of carbon, hydrogen, and oxy- 
gen. Ordinary or ethyl alcohol is the best known member 

Some Common Organic Compounds. 409 

of this group. It is usually called simply alcohol. There 
are many alcohols analogous to ethyl alcohol, but the only 
other important one is methyl alcohol. 

The alcohols may be regarded as hydroxides of certain radicals, e.g. 
ethyl, methyl, propyl, etc. 1 For example, ethyl alcohol is ethyl hydrox- 
ide, and may be considered as formed by replacing one hydrogen atom 
of ethane (C 2 H (i ) by one hydroxyl group (OH). Again, alcohols are 
analogous to metallic hydroxides, in which the metal is replaced by a 

Ethyl Hydroxide Sodium Hydroxide 

Alcohols and metallic hydroxides have some properties in common. 
Thus, both form salts with acids. With acetic acid, sodium hydroxide 
forms sodium acetate, while alcohol forms ethyl acetate (see Ethereal 

Methyl Alcohol, CH 3 .OH, is a colorless or slightly 
yellowish liquid, much like ordinary alcohol. It boils at 
about 66 C, and burns with a pale flame which de- 
posits no soot. It intoxicates, and if concentrated is 
poisonous. It mixes with water in all proportions. It is 
cheaper than ethyl alcohol, and is used as a solvent for 
fats, oils, and shellac, and in the manufacture of varnishes 
and dyestuffs. Methyl alcohol is often called wood alco- 
hol or wood spirit, because it is one of the liquid products 
obtained by the dry distillation of wood (see Charcoal). 

Ethyl Alcohol, C 2 H 5 . OH, is a colorless, volatile liquid, 
having a burning taste and a pleasant odor. It is lighter 
than water, its specific gravity being about 0.8. It boils 
at 78.3 C., and does not freeze until at 130.5 C. Be- 
cause of its very low freezing point, it is used in ther- 

1 The names of these and similar radicals are derived from the correspond- 
ing hydrocarbon. Thus, the word methyl comes from methane, ethyl from 
ethane, propyl from propane. 

4i o Descriptive Chemistry. 

mometers designed to record temperatures below 40 C. 
(the freezing point of mercury), as in Arctic explorations. 
Its harmful effect on the human system need not be dis- 
cussed. Alcohol mixes with water in all proportions. 
The ordinary commercial variety contains from 50 to 95 
per cent of alcohol. Pure or absolute alcohol is obtained 
by removing the remaining water with lime. Proof spirit 
contains about 50 per cent of alcohol. Methylated spirit 
contains 90 per cent ethyl and 10 per cent methyl alcohol; 
it is often used as a cheap substitute for ordinary alcohol, 
but it cannot be used as a beverage on account of the dis- 
agreeable taste imparted by the methyl alcohol. Alcohol 
is an excellent solvent for gums, oils, and resins, and is 
therefore extensively used in the manufacture of varnishes, 
essences, extracts, tinctures, perfumes, and medicines. It 
is also used as an antiseptic, and as a source of heat in 
alcohol lamps. Many organic compounds, as ether and 
chloroform, are prepared from alcohol. Some vinegar is 
made from alcohol. In museums alcohol is used to pre- 
serve specimens. Alcohol may be prepared from ethane (see 
below), but it is manufactured by the fermentation of sugars. 
Fermentation is a general term for the chemical changes 
caused by ferments. The latter are usually minute living 
bodies, though some inorganic chemical sub- 
stances cause fermentation. The process and 
essential products vary with the nature of the 
ferment. The important kinds of fermenta- 
tion are alcoholic, acetic, and lactic, and the 
respective products are alcohol, acetic acid, 
FlG -77- and lactic acid. Alcoholic fermentation is 

Yeast cells. 

caused by ordinary yeast. Under the micro- 
scope, yeast has the form of slimy yellow chains of small, 
round cells (Fig. 77). When yeast is added to a solution 

Some Common Organic Compounds. 411 

of glucose, or any other fermentable sugar, the yeast 
plants multiply rapidly. Air must be admitted, and the 
temperature should be 2O-3O C. The changes are 
numerous and complex, but the main products are alcohol 
and carbon dioxide, thus 

C 6 H 12 6 2C 2 H 6 + 2C0 2 

Glucose Alcohol Carbon Dioxide 

The fermentation ceases as soon as the liquid contains 
about 14 per cent of alcohol. The solution is filtered and 
concentrated by distillation, until the distillate contains the 
desired per cent of alcohol. Commercial alcohol is made 
also from potatoes, grains, rice, beet root, molasses, and 
many other substances rich in sugar and starch. Ordinary 
or cane sugar must be boiled with acid before it will 

Wines, beers, and all alcoholic liquors are prepared by 
fermentation. Yeast is seldom added, however, because 
the ferment which brings about the change is in the air, 
upon fruits and vines. Wines are made from the juice of 
grapes ; beer is made from hops and malt (barley which 
has sprouted). Whisky, gin, brandy, rum, and cordials 
are called distilled liquors, and are manufactured by dis- 
tilling the liquid obtained by fermenting grains, molasses, 
fruit juices, and other substances containing sugar and 
starches. Hence, wine, beer, and similar liquors are essen- 
tially mixtures of alcohol and water. They differ mainly 
in their proportion of alcohol. The particular flavor is due 
to small quantities of different substances which are inten- 
tionally added, obtained from the raw materials, or formed 
by special processes of manufacture. Coloring matter is 
usually added, but sometimes it is extracted from the casks 
in which the liquor is stored. Beer contains from 3 to 7 

412 Descriptive Chemistry. 

per cent of alcohol, wines from 6 to 20, rum, brandy, and 
whisky from 40 to 60 or more per cent. 


Aldehydes are compounds of carbon, hydrogen, and oxy- 
gen. They are formed by the oxidation of alcohols. The 
two important members of this group are acetic aldehyde 
(or acetaldehyde) and formic aldehyde (or formaldehyde). 

Acetic Aldehyde, CH 3 . CHO, is usually called simply aldehyde. It 
is a colorless, very volatile liquid, and has a peculiar, suffocating odor. 
It is a vigorous reducing agent, and is sometimes used to precipitate 
silver, as a thin coating, from silver solutions. It is converted by oxi- 
dizing agents into acetic acid (hence its name, acetic aldehyde}. Alde- 
hyde is prepared by oxidizing alcohol with a solution of potassium (or 
sodium) dichromate and sulphuric acid. When a mixture of these three 
substances is gently warmed, the characteristic odor of aldehyde may be 
detected. The oxidation of alcohol consists simply in the removal of 
hydrogen, thus 

C 2 H,.OH + O = CH 3 .CHO + H 2 O 
Alcohol Aldehyde 

The word aldehyde emphasizes this fact, being a contraction of 0/cohol 

When chlorine is used to oxidize alcohol, part of the hydrogen is 
replaced by chlorine, and the compound CC1 3 . CHO is formed. This 
substance, called chloral, forms a hydrate (CC1 3 . CHO . H 2 O), which is 
used to induce sleep and relieve pain. When chloral is treated with an 
alkali, it is decomposed and chloroform (CHC1 3 ) is produced. The 
latter is a sweet liquid, and is used to produce insensibility in surgical 
operations. Chloroform is usually made by treating alcohol with bleach- 
ing powder. lodoform (CHI 3 ), which is analogous to chloroform, is a 
yellow solid, with a disagreeable smell, and is extensively used as a 
dressing for wounds. It protects the wound from the harmful action of 

Formaldehyde, H . CHO, is a gas, but is used only in 
solution. It has a penetrating odor. The commercial solu- 

Some Common Organic Compounds. 413 

tion sold as formalin contains 40 per cent of formaldehyde. 
It corresponds to methane and methyl alcohol, thus 

H H H 

I I I 

H-C-H H-C-O-H C = O 

I I I 

I H H H 

Methane Methyl Alcohol Formaldehyde 

With oxygen it forms formic acid (hence its name, see 
below). Large quantities of formaldehyde are used in the 
manufacture of dyestuffs and fuming nitric acid, as a food 
preservative, and a disinfectant. When used for the last 
purpose, the solution is vaporized in a special kind of lamp, 
and the vapors are conducted by a small tube into the room 
to be disinfected. It is one of the most convenient and 
efficient of all disinfectants, and is very generally used. 



Ethers are compounds of carbon, hydrogen, and oxygen. 
They are analogous to the metallic oxides. They are 
formed by heating alcohols with sulphuric acid. Ordinary 
or ethyl ether is the best known member of this group. 

Ethyl Ether, C 4 H 10 O, is a colorless, volatile liquid, with 
a peculiar, pleasing taste and odor. It is lighter than 
water, its specific gravity being about 0.74. It boils at 
35 C, and the vapor is very inflammable. The liquid 
should never be brought near a flame. It is somewhat 
soluble in water, and it also dissolves water to a slight 
extent. It mixes with alcohol in all proportions. It is a 
good solvent for waxes, fats, oils, and other organic com- 
pounds. Its chief use is as an anaesthetic, i.e. to render one 
insensible to pain in surgical operations. 

414 Descriptive Chemistry. 

Ether is manufactured by distilling a mixture of ethyl alcohol and 
sulphuric acid in the proper proportions. Hence, the names, ethyl 
or sulphuric ether. Ethylsulphuric acid is first produced, thus 

C 2 H 5 .OH + H 2 S0 4 HC 2 H 5 SO 4 + H 2 O 

Alcohol Sulphuric Acid Ethylsulphuric Acid 

When more alcohol and the ethylsulphuric acid are heated together, 
ether is formed, and sulphuric acid is reproduced, thus, 

HC 2 H 5 SO 4 + C 2 H 5 . OH = (C 2 H 5 ) 2 O + H 2 SO 4 


The process is thus continuous, a small quantity of sulphuric acid serv- 
ing to transform a large quantity of alcohol into ether. Ethyl ether is 
ethyl oxide, (C 2 H 5 ) 2 O or C 2 H 5 . O . C 2 H 5 . 


Organic Acids are compounds of carbon, hydrogen, and 
oxygen. It is a large class of compounds divided into 
several series, one of the most important of which is the 
acetic or fatty series. Its best known member is acetic 
acid ; several of the higher members occur in fats and oils. 

These acids are closely related to hydrocarbons, alcohols, and alde- 
hydes, as may be seen by the following formulas : 

H H 

I I 

H-C-H H-C-(OH) 

I I 

H-C-H H-C-H 

I I 

H H 

Ethane Ethyl Alcohol Acetic Aldehyde Acetic Acid 

It is thus possible to pass from a hydrocarbon through a correspond- 
ing alcohol and aldehyde to an acid. 

The characteristic group of atoms in organic acids is COOH (or 
O = C - O - H), and is called carboxyl. 



C = O 










Some Common Organic Compounds. ,415 

Acetic Acid, C 2 H 4 O 2 or CH 3 . COOH. This is the 
most common organic acid. It is manufactured on a large 
scale by the dry distillation of wood. The dark red 
watery distillate, which is called pyroligneous acid, con- 
tains about 10 per cent of acetic acid besides a small per 
cent of methyl alcohol and many other organic compounds. 
This distillate is neutralized with lime or sodium carbonate, 
and the acetate formed is then decomposed and distilled 
with hydrochloric or sulphuric acid. The acetic acid which 
condenses in the receiver may be further purified by dis- 
tilling it with potassium dichromate and then filtering 
through charcoal. Sometimes the pyroligneous acid is 
distilled without neutralizing ; the distillate is then dilute, 
impure acetic acid, known as wood vinegar. If sodium 
acetate, prepared as described above, is fused and then 
distilled with concentrated sulphuric acid, the product is 
a very concentrated acetic acid. It is called glacial acetic 
acid, because at about 1 7 C. it becomes an icelike solid. 

Commercial acetic acid is a water solution containing 
about 30 per cent of pure acetic acid. It is a colorless 
liquid, having a pleasant odor and a sharp taste. It is 
slightly heavier than water. It mixes with water and alco- 
hol in all proportions, and like alcohol is an excellent 
solvent for many organic substances. Recently, it has 
begun to replace alcohol as a solvent for many drugs. 

Acetic acid is used to prepare acetates, dyestuffs, and 
other organic compounds, medicines, white lead, and in 
the manufacture of vinegar. 

Vinegar is dilute, impure acetic acid. It is prepared by 
oxidizing dilute alcohol, the essential change being repre- 
sented thus 

C 2 H 6 4- 2 = C 2 H 4 2 + H 2 
Alcohol Oxygen Acetic Acid Water 


Descriptive Chemistry. 

The transformation is accomplished by fermentation. 
Two processes are used, (i) When beer, weak wines, or 
cider are exposed to the air, they slowly become sour, 
owing to the conversion of alcohol into acetic acid. The 
change is caused by the presence and activity of a ferment, 
known as mycoderma aceti, or " mother of vinegar." Strong 
wines and pure dilute alcohol do not become sour, because 
the ferment cannot live in such liquids. (2) In the "quick 
vinegar process," impure dilute alcohol is oxidized by ex- 
posing it to an excess of air. The operation is conducted 

in tall vats or casks filled with 
beechwood shavings soaked 
in strong vinegar (Fig. 78). 
Holes at the bottom and top 
allow air to enter and escape 
freely. The alcoholic solu- 
tion is introduced at the top, 
trickles through the shavings, 
and collects at the bottom. 
In its passage it comes in 
contact with the ferment and 
oxygen, and is partially con- 
verted into vinegar. The 
operation is repeated until 
Thus prepared, the vinegar 

lacks the flavor, odor, and color of cider vinegar, but these 
deficiencies are often artificially supplied. 

Vinegar is used chiefly as a condiment for the table and 
in making pickles and similar relishes. 

The constitution of acetic acid has been shown to correspond to the 
formula CH 3 . COOH. Its metallic salts are formed by substituting a 
metallic atom (or group) for the hydrogen of the group COOH. 
radical CH 3 remains unchanged. (See page 170.) 

FIG. 78. Apparatus for the prep 
aration of vinegar from impure, dilute 

the change is complete. 


Some Common Organic Compounds. 417 

Acetates. Acetic acid is a monobasic acid, and forms 
a series of salts the acetates. They are prepared like 
other salts by the interaction of the acid and carbonates, 
hydroxides, metals, etc. The metallic acetates are usually 
crystallized solids, which readily yield acetic acid when 
treated with sulphuric or a similar acid. Most of them 
contain water of crystallization, and most are poisonous. 

Several acetates have useful applications. Sodium acetate, 
NaC 2 H 3 O 2 . 3 H 2 O, is a white crystallized solid, used in preparing 
pure acetic acid, and in the manufacture of dyestuffs. Lead acetate, 
Pb(C 2 H s O 2 ) 2 , is a white crystallized solid, used, in dyeing and in mak- 
ing a yellow pigment. Its sweet taste led to the common name of 
"sugar of lead. 11 Aluminium acetate, A1(C 2 H 3 O 2 ) 3 , is not known in 
the pure state, but an impure solution, known as " red liquor," is exten- 
sively used in dyeing and calico printing. Iron acetates are sold in 
solution as a complex black liquid, known as "iron liquor," which is 
used in dyeing black silks and cottons, and in calico printing (see 
Mordants). A complex copper acetate, 2 Cu(C 2 H 3 O 2 ) 2 + CuO, called 
verdigris, is used in making blue paint. Another complex acetate of 
copper and arsenic is Paris green ; it is used to kill potato bugs and 
other insects which injure vegetation. 

A few other acids in this series are interesting. Butyric acid 
C 4 H 8 O 9 , is the acid which gives the disagreeable odor to rancid butter. 
Stearic acid. C 18 H 3(J O 2 , and Palmitic acid, C 16 Ho 2 O 2 , are found as 
compounds in beef suet, mutton fat, butter, and other fats. Palmitic 
acid is also one of the essential compounds found in palm oil. These 
two acids are white solids, and are used to make stearin candles (see 
Fats, below). 

Other Organic Acids which are important are oxalic, 
lactic, malic, tartaric, and citric. 

Oxalic Acid occurs as a salt in rhubarb and sorrel. It 
is manufactured on a large scale by heating sawdust with 
potassium hydroxide, and treating the residue first with 
lime and then with sulphuric acid. Oxalic acid is a white 
solid, very soluble in water, from which it crystallizes with 

4i 8 Descriptive Chemistry. 

two molecules of water of crystallization (C 2 H 2 O 4 . 2. H 2 O). 
It is very poisonous. It is dibasic and forms several use- 
ful salts. The acid and some of its salts decompose iron 
rust and inks containing iron, and are often used to remove 
such stains from cloth. The acid and its salts are also 
used in dyeing, calico printing, photography, in making 
dyestuffs, and as an ingredient of mixtures for cleaning 
brass and copper. 

Lactic Acid, C 3 H 6 O 3 , occurs in sour milk, being one 
product of the fermentation of the milk sugar. It is a 
thick, sour liquid, and is easily decomposed by heat. 
When sour milk is used in cooking, the " baking soda " 
and lactic acid interact, producing soluble sodium lactate 
and carbon dioxide gas. Lactic acid and its salts are used 
as medicines, in beverages, and as a substitute for more 
expensive acids in dyeing and calico printing. 

Malic acid, C 4 H 6 O 5 , is found free and as salts in apples, pears, cur- 
rants, gooseberries, rhubarb, grapes, and berries of the mountain ash 
tree. It is a white, crystalline solid. 

Tartaric Acid, C 4 H 6 O 6 , occurs as the potassium salt in 
grapes and other fruits. During the fermentation of grape 
juice, impure acid potassium tartrate is deposited in the 
casks. From this argol or crude tartar the acid itself 
is prepared by treating the raw product successively with 
chalk and sulphuric acid. Tartaric acid is a white crystal- 
lized solid, soluble in water and alcohol. It is used in dye- 
ing, and as one ingredient of Seidlitz powders. In these 
and similar powders it serves to decompose the other in- 
gredient which is a carbonate (see Sodium Bicarbonate). 

Tartaric acid is dibasic and forms two classes of salts. Purified 
acid potassium tartrate obtained from argol is commonly known as 
cream of tartar. It is extensively used in the manufacture of baking 
powders. These, as a rule, are essentially mixtures of cream of tartar 

Some Common Organic Compounds. 419 

and sodium bicarbonate, HNaCO 3 . When moistened by dough, the 
baking powder dissolves, the two ingredients interact and liberate car- 
bon dioxide as the main product. This gas bubbles slowly through 
the dough, thereby puffing it up and making it porous (see Sodium 
Bicarbonate). Tartar emetic is a tartrate of potassium and antimony. 
It is used as a medicine and to some extent in dyeing. 

Citric Acid, C(;H 8 O 7 , occurs abundantly in lemons and oranges, and 
in small quantities in currants, gooseberries, and raspberries. It is a 
white, crystallized solid, very soluble in water. The taste is sour, but 
pleasant. The acid and its magnesium salt are used as medicines. The 
acid itself is used in calico printing. Citric acid is tribasic. 


Ethereal Salts or Esters are compounds of carbon, hy- 
drogen, and oxygen closely related to alcohols and organic 
acids. Thus, when ethyl alcohol, acetic acid, and concen- 
trated sulphuric acid are mixed and warmed, ethyl acetate 
is formed. The essential change is represented thus ^ 

C 2 H 5 .OH +CH 3 .COOH = CH 3 .COOC 2 H 5 + H 2 O 

Ethyl Alcohol Acetic Acid Ethyl Acetate Water 

The sulphuric acid serves to absorb the water. Ethyl 
acetate has a pleasant, fruitlike odor, and its formation in 
this way is a simple test for alcohol or acetic acid. Ethyl 
acetate is analogous to sodium acetate, i.e. the organic salt 
contains the radical ethyl while the metallic salt con- 
tains sodium. The fatty acids, as well as those of other 
series, form many ethereal salts of special interest. Some 
occur naturally in fruits and flowers, and in many cases 
give the flavor and fragrance. Others are prepared artifi- 
cially and used as the basis of cheap flavoring extracts, 
perfumery, and beverages. Ethyl butyrate has the taste 
and fragrance of pineapples, amyl acetate of bananas, 
amyl valerate of apples. 

420 Descriptive Chemistry. 


General Relations. Natural fats and oils are essentially 
mixtures of stearin, palmitin, and olein. Beef and mutton fat 
are chiefly stearin, lard is mainly palmitin and olein ; while 
oils, such as olive oil, are largely olein. Stearin and pal- 
mitin are solids at the ordinary temperature, but olein is a 
liquid. These three compounds stearin, palmitin, and 
olein are ethereal salts of their corresponding acids and 
the alcohol, glycerine. They are analogous to ethyl acetate. 
The radical of glycerine is glyceryl, C 3 H 5 . Thus, stearin 
is glyceryl stearate, palmitin is glyceryl palmitate, and 
olein is glyceryl oleate. Natural fats and oils, therefore, 
are mixtures of these and similar ethereal salts. Fats are 
sometimes called glycerides. Glycerine is a triacid alcohol 
containing three hydroxyl (OH) groups. Like ordinary 
alcohol, it interacts with the fatty acids and forms ethereal 
salts. The latter, as we have just learned, are the fats. 
Now when fats are heated with very hot steam or with sul- 
phuric acid, the fats themselves are changed into glycerine 
and the corresponding acids. Thus, with stearin, the 
change is 

(C 17 H 35 . C0 2 ) 3 C 3 H 5 + 3 H 2 = C 3 H 5 (OH), + 3 C ir H,, . COOH 

Stearm Glycerine Stearic Acid 

But if fats are boiled with sodium hydroxide or a simi- 
lar alkali, glycerine and an alkaline salt of the correspond- 
ing acid are formed. Soap is a mixture of such alkaline 
salts. In a few words, the general relations are these: 
(i) fats are ethereal salts. (2) Treated with steam or acid, 
fats form glycerine and fatty acids. (3) Treated with alka- 
lies, fats form glycerine and soap. 

Natural Fats and Oils are often complicated mixtures. 
The solid fats, as already stated, are rich in stearin and 

Some Common Organic Compounds. 421 

palmitin. Tallow is chiefly stearin, but human fat and 
palm oil are largely palmitin. The soft and liquid fats and 
oils contain considerable olein, as a rule. The proportion 
of olein determines the consistency of the fats and oils. 
Thus, Olive oil contains about 72 per cent of olein (and a 
similar fat) and 28 per cent of stearin and palmitin. The 
specific character of many fats and oils is due mainly to 
the presence of a small proportion of certain fats. These 
fats correspond to uncommon acids in the fatty, oleic, and 
other series. Butter, for example, consists mainly of the 
fats corresponding to the following acids : palmitic, stearic, 
oleic, butyric, capric, and caproic. The last three with 
traces of other substances give butter its pleasant flavor. 
Oleomargarine and other substitutes for butter resemble 
real butter very closely in composition. Artificial butter, 
however, lacks the flavor of the real butter, but it is " prob- 
ably just as nutritious, although perhaps not quite so easily 
digested." The lack of flavor noticed in artificial butter is 
due to the absence of the fats corresponding to the acids 
of low molecular weight. Cottolene is a mixture of beef 
fat and cotton-seed oil ; it is used as a substitute for lard. 

Glycerine (C 3 H 8 O 3 or C 3 H 5 .(OH) 3 ) is a thick, sweet 
liquid. It mixes readily with water and with alcohol in all 
proportions, and absorbs moisture from the air. Heated 
in the air, it decomposes and gives off irritating gases, like 
those produced by burning fat. 

Glycerine is used to make nitroglycerine (see below), 
toilet soaps, printers' ink rolls ; it is also used as a solvent, 
a lubricator, a preservative for tobacco and certain foods, 
a sweetening substance in certain liquors, preserves, and 
candy ; as a cosmetic ; and, owing to its non-volatile and 
non-drying properties, it is used as an ingredient of inks 
and oils. 

422 Descriptive Chemistry. 

Glycerine is a by-product in the manufacture of soap, or it is made 
directly by decomposing fats with steam under pressure or with lime. 
Ail these methods involve the chemical change described above, viz. 
the decomposition of an ethereal salt (the fat) into the corresponding 
alcohol (glycerine) and a mixture of fatty acids. By skillful treatment 
the glycerine is freed from water and impurities. The mixture of fatty 
acids is made into the so-called "stearin" candles. 

As already stated, glycerine is an alcohol, and for this reason it is 
often called glycerol. When treated with a mixture of concentrated 
nitric and sulphuric acids, it forms an ethereal salt commonly known as 
nitroglycerine (C 3 H 3 (ONO 2 ) 3 ). This is a yellow, heavy, oily liquid. 
It is the well-known explosive, and is also an ingredient of some other 
explosives. When kindled by a flame, it burns without explosion ; but 
if struck by a hammer or heated suddenly by a percussion cap, it ex- 
plodes violently. Nitroglycerine is used in blasting ; but since it is dan- 
gerous to handle and transport, it is usually mixed with some porous 
substance, such as infusorial earth, fine sand, or even sawdust. In this 
form it is called dynamite. 

Soap, as already stated, is a mixture of alkaline salts of 
organic acids, mainly stearic and palmitic acids. Soap is 
made by boiling fats with sodium hydroxide or potassium 
hydroxide. This process is called saponification. Sodium 
hydroxide produces hard soap, consisting chiefly of sodium 
palmitate, sodium stearate, and sodium oleate. Potassium 
hydroxide produces soft soap, which is mainly the corre- 
sponding potassium salts. The chemical change, as already 
stated, consists in tr e transformation of an ethereal salt 
(fat) into glycerine and an alkaline salt. In the case of 
pure stearin (glyceryl stearate) the change may be repre- 
sented thus 

C 3 H 5 (C 17 H 35 . C0 2 ) 3 + 3NaOH - 3 C 17 H 35 . CO 2 Na + C 3 H 5 (OH) 3 
Stearin Sodium Sodium Glycerine 

Hydroxide Stearate 

The fats used in soap making vary with the soap. Tal- 
low, lard, palm oil, and cocoanut oil make white soaps. 

Some Common Organic Compounds. 423 

Bone grease or house grease, together with tallow, palm 
oil, cotton-seed oil, and rosin, make yellow soaps. Olive 
oil is used for making castile soap. 

In the^cold process the calculated amounts of alkali and fat are allowed 
to interact, first in a large tank and then in a box called a " frame." By 
this process the glycerine and excess of alkali are left in the soap. Most 
soaps are made by the boiling process. The fat and alkali are boiled 
in a huge kettle. This operation produces a thick, frothy mixture of 
soap, glycerine, and alkali. At the proper time, salt is added, thereby 
causing the soap to separate and rise to the top. The liquid beneath is 
drawn off, and from it glycerine is extracted. The soap is often boiled 
again with rosin or cocoanut oil ; then purified by washing, mixed, if 
desired, with perfume, coloring matter, or some filling material (such as 
sodium silicate, sand, borax), cooled in "frames," cut, and dried. Most 
soaps contain water. This really assists their cleansing action. The 
latter is believed to be due to the free alkali formed by the decomposi- 
tion of the soap when dissolved. 


Carbohydrates are compounds of carbon, hydrogen, and 
oxygen. This is a large group, and the most important 
members are the sugars, starches, and cellulose. 

The term carbohydrate is applied to these compounds because they 
contain hydrogen and oxygen in the proportion to form water. They 
were once regarded as hydrates of carbon, or carbon hydrates a view 
which is incorrect and misleading. 

Sugars. The popular term sugar means almost any 
sweet substance found in fruits, nuts, vegetables, sap of 
trees, etc., though it is usually restricted to the ordinary 
white sugar obtained from sugar cane and sugar beet. 
Chemically, there are many sugars, each having a defi- 
nite constitution. The most important is ordinary sugar, 
which is also called cane sugar, sucrose, and saccharose. 
Another important sugar is glucose. 

424 Descriptive Chemistry. 

Cane Sugar, C 12 H 22 O n , is widely distributed in nature, 
being found in the sugar cane, sugar beet, sugar maple, 
Indian corn, sorghum, most sweet fruits, many nuts, blos- 
soms of flowers, and honey. The main source of cane 
sugar is the sugar cane and sugar beet. 

Saccharose, or ordinary sugar, is a white, crystallized 
solid. Rock candy is highly crystallized sugar. It is solu- 
ble in water, but only sparingly soluble in alcohol. Heated 
to 160 C, sugar melts, and on cooling forms a pale yellow 
colored mass, called barley sugar. Heated to about 200 C., 
it is changed into water and a brown mass, called caramel, 
which is used to color liquors, soups, etc. If sugar is 
heated with sulphuric acid, it is changed into a black mass, 
which is mainly carbon ; several gases are also produced, 
such as steam, carbon dioxide, and sulphur dioxide. Cane 
sugar does not ferment. 

The manufacture of Cane Sugar from sugar cane and sugar beets 
involves two main operations: (i) the preparation of raw sugar and 
(2) its purification or refining, (i) In the preparation of raw sugar 
from sugar cane the juice is extracted from the cane by crushing the 
latter between heavy iron rollers. The liquid is then clarified as soon 
as possible by boiling it with a little lime, removing the scum which 
contains much of the impurity, and finally filtering the liquid through 
bags or a filter press. The purified juice is next evaporated until the 
cane sugar begins to crystallize from the cooled liquid. Formerly the 
evaporation was accomplished in an open pan, and is now in some 
localities, but usually a vacuum kettle is used. The crystals are next 
separated from the liquid by allowing the latter to drip out, or more 
commonly by whirling it out in a centrifugal machine. The solid 
product is called muscovado, raw or brown sugar. The thick liquid 
is the familiar molasses. There are several grades of each product. 
The preparation of raw sugar from sugar beets resembles the method 
used for sugar cane. The washed beets are reduced to a pulp, or cut 
into slices, and then treated with water. The sugar dissolves in the 
water. The solution is clarified, evaporated, and separated by pro- 
cesses much like those applied to cane-sugar solutions. The raw sugar 

Some Common Organic Compounds. 425 

can scarcely be distinguished from cane sugar. The molasses is unfit 
for table use, though considerable sugar is extracted from it by means 
of strontium hydroxide (see Strontium Hydroxide). (2) Raw sugar is 
usually dark colored, and must be refined before it is suitable for most 
uses. The refining of sugar consists in (a) purification, and (<) recrys- 
tallization. () The raw sugar is purified by first dissolving it in huge 
tanks. Air is blown in to agitate the heated solution, blood and other 
substances are often added to entangle the impurities, and lime is also 
added to precipitate and gather the impurities into a scum or clot. 
The colored liquid is next filtered, first through cloth bags and then 
through animal charcoal, from which it drips as a perfectly clear liquid, 
(b} The filtered sirup is now evaporated in a large vacuum kettle. 
When a sample shows that the evaporation has reached the proper 
point, the liquid is run into tanks to crystallize. The crystals of sugar 
are separated from the sirup by centrifugal machines. The latter is 
boiled again or sold as sirup for the table. The crystals are dried in a 
heated tube called a granulator, so that each grain will be separate. 
Hence the name granulated sugar. The grains are sifted and packed 
in barrels for the market. 

Lactose, or sugar of milk, has the same formula as cane sugar, but 
its constitution and properties differ. It is obtained from milk. Its 
crystals are white, hard, gritty, less sweet than cane sugar ; they con- 
tain one molecule of water of crystallization. Sugar of milk is used in 
making homeopathic pills and certain kinds of foods for infants. 

Glucose is the name of a sugar and of a commercial 
mixture of glucose and several related substances. Glu- 
cose (dextrose or grape sugar, C 6 H 12 O 6 ) is found in many 
sweet fruits, especially in grapes. Old raisins are some- 
times coated with this sugar. It is often associated with 
levulose (fructose or fruit sugar) an is6meric compound 
(C 6 H 12 O 6 ). The two sugars are found, for example, in 
honey and in parts of some plants. Both sugars are 
formed from cane sugar by boiling it with a dilute acid. 
The chemical change may be represented thus 

C 12 H 22 O n + H 2 O = C 6 Hi 2 O 6 + C 6 H 12 O 6 
Cane Sugar Glucose Fructose 

426 Descriptive Chemistry. 

Both glucose and fructose ferment, forming alcohol and 
carbon dioxide (see Alcohol). 

The commercial mixture called "glucose" is prepared on a large 
scale by boiling starch with a dilute acid, usually sulphuric acid. The 
consistency and composition of the product vary with the details of 
manufacture. The liquid products are called " glucose " or "mixing 
sirup," while the solid product is known as " grape sugar " or " dex- 
trose." All contain more or less glucose and are about three fifths as 
sweet as sugar. But since they dissolve in water, and are cheaper than 
cane sugar, they are used extensively in the manufacture of candy, jelly, 
table sirups, etc. They are also added to wines and liquors, certain 
medicines, and many thick liquids in which their presence is harmless. 
In alkaline solutions, glucose is a strong reducing agent, and is used as 
such in dyeing with indigo. It also reduces an alkaline mixture of cop- 
per sulphate, known as Fehling's solution. When this solution is 
boiled with glucose, a reddish copper compound (cuprous oxide) is 
formed. The presence of sugar in solution is often shown in this way. 

Starch is widely distributed in the vegetable kingdom. 
It is found in wheat, corn, and all other grains, in pota- 
toes, beans, peas, and similar vegetables, and in large 
quantities in rice, sago, tapioca, and nuts. Many parts of 
plants contain starch, for example, the stalk, stem, leaves, 
root, seed, and fruit. The food value of vegetables de- 
pends largely upon the starch they contain. 

FIG. 79. Starch grains (magnified) wheat (left), rice (center), corn (right). 

Starch is a white powder, as usually seen. But under 
the microscope it is found to consist of a mass of oval 

Some Common Organic Compounds. 427 

grains, varying somewhat with the source (Fig. 79). 
Starch is only very slightly soluble in water. But if 
heated with water, the grains swell and burst, partially 
dissolve, and form a solution which, when cold, becomes 
the familiar starch paste. Starch in solution is turned 
blue by iodine, and its presence in many vegetables and 
foods may be readily shown by grinding the substance in 
a mortar with warm water and adding a drop of iodine 

Starch is prepared on a large scale chiefly from corn and potatoes. 
The operation is mainly mechanical, and consists in separating the 
starch from the fatty, nitrogenous, and mineral matters in the raw 
product. Immense quantities are consumed as food, in laundries, in 
finishing cloth and paper, in making glucose, and as a paste. 

The composition of starch, according to some authorities, corre- 
sponds to the formula C 6 H 10 O,, but its formula is still being investigated. 

Dextrin is a sticky solid formed from starch by heating 
it to 2OO-25O C. or by treating it with dilute acids. It 
is soluble in water and forms a sticky solution. Commer- 
cial dextrin or British gum is a mixture of dextrin and 
similar compounds. Mucilage contains dextrin. Large 
quantities are used as the gum for the backs of postage 
stamps, and for sticking the colors to the cloth in calico 

Dextrin is sometimes regarded as an intermediate product between 
starch and dextrose. Its composition, according to some authorities, 
corresponds to the formula C 12 H 20 O 10 , but the statement made about the 
composition of starch also applies to dextrin. 

Bread. Wheat flour contains about 70 per cent of starch. The re- 
mainder is chiefly water and gluten in nearly equal proportions, though 
small quantities of mineral matter, dextrin, and other fermentable sub- 
stances are present. In making bread, flour, milk or water, and a little 
yeast are thoroughly mixed into dough, which is put in a warm place to 
rise, Fermentation begins at once. The yeast changes the ferment- 

428 Descriptive Chemistry. 

able substances into alcohol and carbon dioxide. The gases, in trying 
to escape, puff up the dough, which literally rises and becomes light and 
porous. When the dough is baked, the heat kills the yeast, and fer- 
mentation stops ; but the alcohol, carbon dioxide, and some water escape 
and puff up the mass still more. The heat, however, soon hardens the 
starch, gluten, etc., into a firm but porous loaf. 

Cellulose (C 6 H 10 O 5 ) n is widely distributed in the vegetable 
kingdom. The framework of all vegetables is cellulose. It 
is thus analogous to the bones of animals. Wood, cotton, 
linen, and paper are largely cellulose. Pure cellulose is a 
white substance, insoluble in most liquids, but soluble in a 
mixture of ammonia and copper oxide. Concentrated sul- 
phuric acid dissolves it slowly ; and if the solution is di- 
luted and boiled, the cellulose is changed into a mixture 
of glucose and dextrin. By this operation, wood could be 
made into a sugar and then into alcohol ; but the method 
would be too expensive to use on a large scale. 

Sulphuric acid of a certain strength, if quickly and properly applied to 
paper, changes it into a tougher form called parchment paper. The 
latter is often substituted for animal parchment (e-g. sheepskin), and 
has a variety of uses. 

Cellulose has properties resembling those of alcohol. Thus it inter- 
acts with acids and forms ethereal salts. With nitric acid it forms cellu- 
lose nitrates, just as glycerine forms glycerine nitrates (see Nitroglyce- 
rine). The cellulose nitrates are the basis of smokeless gunpowders. 
One of the cellulose nitrates is gun cotton. It looks like ordinary cotton, 
and may be spun, woven, and pressed into cakes. It burns with a large 
flame if unconfined ; but when ignited by a percussion cap or when 
burned in a confined space, gun cotton explodes violently- It is used in 
blasting. Other cellulose nitrates are known. Their solution in a mix- 
ture of alcohol and ether is called collodion. When poured or brushed 
upon a glass plate or the skin, the solvent evaporates, leaving behind a 
thin film. It is used in preparing certain photographic material and as 
a coating for wounds. The " new skin " liquid recently offered for sale 
is mainly collodion. It protects wounds from dusty, impure air, and 
thereby facilitates the healing. A mixture of camphor and cellulose ni- 

Some Common Organic Compounds. 429 

trates is called celluloid. It is easily molded into various shapes. The 
white celluloid is made into collar buttons, and the colored varieties are 
made into toilet articles and ornaments. Celluloid smells of camphor, 
can be lighted with a match, and burns freely with a smoky flame. 

Paper is chiefly cellulose. Formerly it was made from 
various kinds of rags ; but now it is made almost entirely 
from wood, especially the paper used for newspapers and 
cheap books. The best paper, such as writing paper, is 
still made from linen rags. 

In making paper from wood, the latter is reduced to a pulp, which 
is washed, spread on a frame or an endless wire gauze, dried, and 
pressed. The pulp is prepared by two processes, the mechanical and 
the chemical. Mechanical pulp is made by holding a stick of wood 
against revolving stone upon which water constantly falls. Chemical 
pulp is made by heating chipped wood with caustic soda, or with cal- 
cium acid sulphite (usually called bisulphite). The operation is con- 
ducted under pressure in huge tanks called digesters. Chemical pulp 
has longer and stronger fibers than mechanical pulp. The two kinds of 
pulp are often mixed. Most paper is loaded, that is, clay, gypsum, or 
other mineral matter is mixed with the pulp to give the paper body. 
Paper intended for printing or writing is sized, that is, the surface is 
coated with gelatine, rosin, or a similar substance to prevent the ink 
from spreading. Many kinds are also smoothed by passing them 
between heavy rollers. Blotting and tissue papers are not sized or 


Introduction. The hydrocarbon benzene was mentioned 
in Chapter XV as the first member of an homologous series. 
In the same chapter coal tar was described as a black, 
complex liquid obtained as a by-product in the manufacture 
of illuminating gas. Now, coal tar is the chief source of 
benzene and some of its related compounds, while from 
benzene itself hundreds of derivatives have been prepared. 
Some are absolutely indispensable to man, but many have 

430 Descriptive Chemistry. 

as yet merely scientific interest. Only the most important 
benzene compounds can be described in this book. 

Benzene, C 6 H 6 , is a colorless liquid, lighter than water, 
and has an odor suggesting coal gas. It burns with a 
luminous, smoky flame, owing to its richness in carbon. 
Ordinary illuminating gas owes its luminosity partly to 
benzene. It dissolves fats, resins, iodine, sulphur, and 
rubber. Benzene is sometimes called benzol. It should 
not be confused with benzine, which is a mixture of hydro- 
carbons derived from petroleum. Benzene is chiefly used 
in preparing its derivatives. 

The Constitution of Benzene has been carefully studied. For rea- 
sons too extended to state here, it is believed that in a molecule of 
benzene the carbon atoms are arranged in a ring. The structural for- 
mula is often written thus 


/ \ 
H-C C-H 


H-C C-H 

\ ^ 


Benzene forms many derivatives. In all of them the six carbon 
atoms remain as a nucleus. No carbon atom can be removed from the 
benzene molecule without producing complete decomposition. But for 
the six hydrogen atoms, other atoms or radicals can be substituted. 
Hence, the almost infinite number of derivatives of benzene. 

Toluene, C H~ . CH 3 , is the second member of the benzene series. 
It may be regarded as methyl benzene ; or as phenyl methane, that is, 
methane (CH 4 ) in which one hydrogen atom is replaced by the radical 
phenyl (C 6 H 5 ). Toluene is obtained from coal tar, and resembles 
benzene in its properties. 

Nitrobenzene, C,.H 5 . NO 2 , is a yellow liquid formed by the inter- 
action of benzene and nitric acid. It is volatile, and has the odor of 

Some Common Organic Compounds. 431 

bitter almonds. Although poisonous, it is used to produce the flavor 
of almonds in essences and perfumery. It is chiefly used, however, in 
the manufacture of aniline. 

Aniline, C 6 H 5 .NH 2 , is an oily liquid, slightly heavier 
than water. It is prepared on a large scale by reducing 
nitrobenzene with nascent hydrogen. From aniline are 
made many compounds known as aniline dyes. The 
starting point of these dyes is rosaniline, which is pre- 
pared by oxidizing a mixture of aniline and toluidine 
(C 6 H 4 . CH 3 . NH 2 ). Derivatives of rosaniline produce 
exceedingly brilliant colors in every variety of shade. 
Vast dyeing industries have risen since the value of coal 
tar was discovered (about 1860). 

Phenol, C 6 H 5 .OH, is a white crystalline solid. It has 
a smoky odor, is poisonous, and burns the skin. Coal tar 
is the source of phenol. A solution of phenol in water, 
popularly called carbolic acid, is used as a disinfectant. 

Derivatives of Phenol are important. Picric acid, or trinitrophenol 
(C (i H 2 (NO 2 )oOH), is a yellow crystalline solid used in dyeing silk yellow. 
Salts of picric acid the picrates are used in making explosives. 
Related to phenol are hydroquinone (C 6 H 4 (OH) 2 ) and pyrogallic 
acid (C (; Ho(OH) 3 ), which are used extensively as developers in 

Acids, Aldehydes, and Ethereal Salts of the Benzene 
Series. The simplest acid is benzoic acid (C 6 H 5 . COOH). 
It occurs in certain balsams and gums. It is usually pre- 
pared from gum benzoin, and is a white crystalline solid 
with a fragrant odor. The corresponding aldehyde (ben- 
zoic aldehyde, C 6 H 5 .COH) is commonly called oil of 
bitter almonds. It is a fragrant liquid and is used to 
some extent as a flavoring substance. Salicylic acid 
(C 6 H 4 . OH. COOH) is a white crystalline solid, which is 
extensively used as a food preservative. Sodium salicylate 

432 Descriptive Chemistry. 

is a common remedy for rheumatism. The corresponding 
aldehyde gives the fragrance to the wild flower known 
as meadowsweet; and methyl salycilate is the essential 
ingredient of the checkerberry. 

Naphthalene, C 10 H 8 , is a white, lustrous, crystalline 
solid obtained from coal tar. It has a penetrating, un- 
pleasant odor, and is used as a substitute for camphor 
under the name of " moth balls." Large quantities of 
naphthalene are used in making dyestuffs. 

Anthracene, C 14 H 10 , is a white crystallized solid, and, 
like naphthalene, is obtained from coal tar. It is one of 
the most important hydrocarbons, because from it alizarin 
is made. Alizarin is a valuable dyestuff, not only because 
it produces brilliant colors with different mordants, but 
also because most of these colors are fast, that is, they 
do not fade like many aniline colors. The Turkey red -so 
common on cotton goods, is produced by alizarin. Aliza- 
rin was formerly obtained from madder root, but now vast 
quantities are artificially prepared. 

Glucosides are substances occurring in many plants and vegetables. 
By the action of ferments they are changed into glucose and other 
substances that are benzene derivatives. Amygdalin, for example, is 
found in bitter almonds, cherry and peach kernels, and laurel leaves. 
The ferment emulsin, which also occurs in the plants, breaks up the 
amygdalin into oil of bitter almonds, hydrocyanic acid, and glucose. 
Tannin is also a glucoside. The tannins are a group of related com- 
pounds found in the leaves, bark, and other parts of the oak, hemlock, 
and pine trees, in sumach, gallnuts, tea, coffee, and numerous plants. 
Several acids have been obtained from tannins. The best known are 
gallic acid and tannic acid ; the latter is also often called simply tan- 
nin, and probably all tannins contain some tannic acid. Tannic acid 
changes into gallic acid according to the following equation 

C 14 H 10 9 + H 2 2C 7 H fi 5 

Tannic Acid Gallic Acid 

Some Common Organic Compounds. 433 

The formula of gallic acid may be written C (; H 2 (OH) 3 . COOH, thus 
showing its relation to benzene. Tannin, in whatever form, produces 
black compounds with iron salts. Its presence in tea, hemlock bark, 
etc., may be shown by the formation of a black precipitate upon the 
addition of ferrous sulphate. This property is utilized in making 
writing ink, though some kinds of ink are now made from aniline 
dyes. The tannin in oak and hemlock barks is used in tanning leather. 
When raw hides are soaked in solutions of tannin, the tannic acid 
changes certain substances in the skin into insoluble compounds, 
which remain in the hide, thereby converting it into the soft pliable 
form known as leather. Tannins are also used as mordants in dyeing 
silk, cotton, and linen. 

Alkaloids are complex compounds obtained from plants and vegeta- 
bles. The chief property is the power to produce marked physiological 
effects upon animals. All of them contain nitrogen, and resemble 
ammonia in having an alkaline reaction and in uniting directly with 
acids to form salts. Their commercial form is usually a salt. Many 
are used as medicines and drugs, although they are poisonous, especially 
if taken in large quantities. Theine or caffeine is the alkaloid obtained 
from tea and coffee. Nicotine comes from tobacco and is very poison- 
ous. Cocaine is obtained from the coca plant. One of its salts is 
used by surgeons and dentists to relieve pain. Quinine and cinchonine 
are extracted from the bark of the cinchona tree ; both are used as a 
remedy for fevers. Morphine is the chief alkaloid found in opium. The 
latter is the dried sap obtained from a certain part of the unripe poppy. 
Morphine in different forms is used to relieve pain and induce sleep. 
The two familiar medicines, laudanum and paregoric, contain prepara- 
tions of opium. Large doses of any form of opium may be fatal. 


1. How were organic and inorganic compounds' once defined ? Do 
they differ fundamentally ? What compounds are now included by the 
term organic? 

2. What is the essential element in organic compounds ? What 
other elements are often present ? 

3. Give four reasons for the vast number of organic compounds. 

4. Define an organic radical. Name three. 

5. Define constitution. Illustrate it by the empirical, rational, and 
graphic formulas of alcohol. 

434 Descriptive Chemistry. 

6. Name the nine important groups of organic compounds. 

7. Review the general properties of hydrocarbons (see Chapter XV) . 
Name four hydrocarbons. 

8. Define an alcohol. Discuss the constitution of alcohols. 

9. Describe the preparation of methyl alcohol. State its properties 
and uses. Why is it called (a} methyl alcohol, and (b) wood alcohol ? 

10. State (a) the properties, and (b) the uses of ethyl alcohol. 
n. What is (a) alcohol, (b) ethyl alcohol, (c) absolute alcohol, 
(d) methylated spirit, (e) proof spirit ? 

12. What is fermentation ? What are ferments ? 

13. Describe the preparation of alcohol. Discuss the preparation, 
composition, and properties of () wines and beers, and (<) distilled 

14. What are aldehydes ? How are they related to alcohols and to 
hydrocarbons ? 

15. Describe the preparation and properties of (a} acetic aldehyde, 
and (b) formic aldehyde. State the uses of the latter. What is its 
commercial name ? 

1 6. What are ethers ? How are they related to alcohols ? 

17. Describe the preparation, and state the properties and uses of 
ordinary ether. 

1 8. What are organic acids ? Illustrate (by acetic acid) their rela- 
tion to hydrocarbons, alcohols, and aldehydes. 

19. Describe the manufacture of acetic acid. State (#) its properties, 
and () its uses. 

20. What is (a) pyroligneous acid, () glacial acetic acid, (<:) wood 
vinegar, (d) commercial acetic acid ? 

21. Discuss the composition of acetic acid. 

22. What is vinegar ? Describe its manufacture. State its proper- 
ties and uses. 

23. What are acetates ? State their general properties. Describe 
four, and state their uses. 

24. Name three other acids (besides acetic) in the fatty acid series. 
Why is this series so called ? 

25. State the occurrence, properties, and uses of (a*) oxalic acid, 
() lactic acid, (c) tartaric acid, (d} citric acid. Where is malic acid 
found ? 

26. What is (a} argol, (b) crude tartar, (c) cream of tartar, (d) tar- 
tar emetic ? 

Some Common Organic Compounds. 435 

27. Review baking powder (see Sodium Bicarbonate). 

28. What are ethereal salts ? How are they formed ? Where are 
they found ? Describe ethyl acetate. Name three other ethereal salts 
and state their properties. 

29. What is the test for (a) alcohol, and () acetic acid ? 

30. State clearly the general relations of fats to glycerine and soap. 

31 . Name the chief ingredients of fats and oils. What is (a) tallow, 
(b} butter, (c} oleomargarine, (d) stearin ? 

32. Describe the preparation of glycerine. State its properties and 

33. Discuss the constitution of glycerine. State the properties and 
uses of (a) nitroglycerine, and () dynamite. 

34. What is soap ? Describe its general method of manufacture. 
What is the chemistry of its manufacture ? What fats and alkalies are 
used in making soap ? Describe (a) the cold process, and (b) the boil- 
ing process of soap making. 

35. What are carbohydrates? Why is this term used? Name 
several carbohydrates. 

36. What are sugars ? Name several. 

37. Discuss the distribution of cane sugar. State its properties. 
What is (a) cane sugar, (b) sucrose, (c} saccharose, (d) barley sugar, 
(e) caramel ? For what is the last used ? 

38. Describe the preparation of raw sugar from (a) sugar cane, and 
(b) sugar beets. 

39. Describe the refining of sugar. 

40. What is (a) granulated sugar, {b} brown sugar, (c} molasses ? 

41. What is the sugar of milk ? What is its scientific name ? For 
what is it used ? 

42. What is the formula of glucose ? W T hat other names has glu- 
cose ? Where is glucose found ? What sugar is closely related to 
glucose ? How is glucose formed from cane sugar ? State the equation 
for the reaction. 

43. How is commercial glucose prepared ? What is (a) commercial 
grape sugar, and (b} " glucose " ? State the properties and uses of 
commercial glucose. 

44. Describe the test for sugar. 

45. Discuss the distribution of starch. Describe starch. State its 
properties. What is the test for starch ? 

46. How is starch prepared ? State its uses. 

436 Descriptive Chemistry. 

47. What is the simplest formula of starch ? How does it differ 
from the formula of () cane sugar, and ($) glucose ? 

48. What is dextrin ? How is it prepared ? For what is it used ? 

49. Discuss the chemistry of bread making. 

50. What is cellulose ? Describe pure cellulose. State its properties. 

51. What is (a) parchment paper, () gun cotton, (c) collodion ? 

52. What is the chief constituent of paper ? Describe the manufac- 
ture of paper. 

53. State the source of benzene. State its properties. What is (a) 
benzol, and (b) benzine ? 

54. To what class of organic compounds does benzene belong? 
Why is it such an important compound ? 

55. What is the chemical relation of benzene to (a) toluene, () 
nitrobenzene, (c) aniline, (d} phenol, (e) benzole acid ? 

56. Describe nitrobenzene. What is its chief use ? 

57. Describe aniline. How is it prepared ? For what is it used ? 

58. Describe phenol. What is its source and use ? What is its 
common name ? 

59. State briefly the relation of phenol to (a) picric acid, (]) pi- 
crates, (c} hydroquinone, (//) pyrogallic acid. What is the use of each ? 

60. Describe briefly benzoic acid and benzoic aldehyde. 

61. Describe salicylic acid. State the use of this acid. 

62. Describe naphthalene. What is its popular name ? State its uses. 

63. Describe anthracene. State its use. What is alizarin ? 

64. What are glucosides ? Discuss (a} the occurrence, () the prop- 
erties, and (c) the uses of tannin. What is (a) ink, and () leather ? 

65. What are alkaloids ? Name six. What is their chief property ? 


1. Alcohol is 0.8 as heavy as water. What is the weight of 1200 cc. 
of alcohol ? 

2. If 10 gm. of pure alcohol are burned, what weight of each product 
is formed ? (Equation is C 2 H (; O + 30, = 2 CO 2 + 3 H 2 O.) 

3. Calculate the percentage composition of (#) alcohol (C 2 H 6 O), () 
acetic acid (C 2 H 4 O 2 , (c} cane sugar (C ]2 H 2 ,O n ). 

4. Calculate the simplest formulas of the substances having the com- 
position : (a} carbon = 40, hydrogen = 6.67, oxygen = 53.33 ; () 
carbon = 15.8, hydrogen = 5.26, nitrogen = 36.84, sulphur = 42.1 ; (c) 
carbon = 54.55, hydrogen = 9.09, oxygen = 36.36. 


1. The Metric System. The fundamental unit of this system of 
weights and measures is the meter. It is the unit of length, and is 
39.37 inches long. 

The meter and the other units have multiples and submultiples, 
which are designated by prefixes attached to the particular unit. The 
multiple prefixes are deca-, hecto-, and kilo-, equivalent respectively to 
10, 100, and 1000. The submultiple prefixes are deci-, centi-, and milli-, 
which correspond respectively to o.i, o.oi, and o.ooi. 

The unit of weight is the gram. It is derived from the kilogram, 
which is the weight of a cubic decimeter of water at 4 C. A kilogram 
weighs about 2.2 pounds. Small weights are expressed in terms of the 
gram. Thus, the weight of an object weighing 2 grams, 2 centi- 
grams, and 5 milligrams is 2.025 grams. 

The unit of volume is the liter. It is equal to the capacity of the 
vessel containing a kilogram of water. A liter equals about one quart. 

The relation between the units, multiples, and submultiples is shown 
in the 


































From this table it is evident that 10 milligrams equal I centigram, 10 
centigrams equal I decigram, 10 decigrams equal i gram, and so on. 



Descriptive Chemistry. 

The relation of the metric system to weights and measures in com- 
mon use is shown by the 


meter = 39.37 inches 
kilometer = 0.62 mile 
centimeter = 0.39 inch 
liter = 0.908 quart 

liter = 1.056 quart '(liq.) 

gram = 15.432 grains 

kilogram = 2.2 pounds (avoir.) 
metric ton = 2204 pounds 


cubic inch 
quart (liq.) 
pound (avoir.) 
ounce (avoir.) 
ounce (troy) 
grain (apoth.) 

2.54 centimeters 

1.6 kilometers 

16.39 cubic centimeters 

0.9465 liter 

0.4536 kilogram 

28.35 grams 

31.1 grams 

0.0648 gram 

The passage from the English to the metric system may be accom- 
plished by utilizing the 




Inches to centimeters 
Centimeters to inches 


Cubic centimeters to cubic inches 


Grams to ounces (avoir.) 


Grams to grains 

J 543 

The customary abbreviations of the common denominations are 

meter, m. 
decimeter, dm. 
centimeter, cm. 

liter, 1. 

kilogram, kg. or Kg. 

decigram, dg. 

cubic centimeter, cc. 
milligram, mg. 
centigram, eg. 

The ^referable abbreviation for gram is gm. The same abbreviation 
is used for singular and plural, e.g. I m., 4 gm., 3 cm., 50 cc. 

A convenient relation (true only in the case of water) to remember 
is i I = i kg. = i cu. dm. = 1000 cc- ^ 1000 gm. = 2,3 lb. 




1. What is the abbreviation of gram, centigram, liter, meter, cubic 
centimeter, centimeter, decimeter, milligram ? 

2. Express (a) i liter in cubic centimeters, () 2 1. in cc., (c) i meter 
in centimeters, (d) 250 cm. in dm., (e) i kg. in grams, (/) 250 gm. 
in mg. 

3. Add 2 kg., 5 dg., 2 eg., 4 gm., and 7 mg., and express the sum in 

4. How many cc. in a liter ? 

5. What is the weight in grams of (a) i liter of water, (<) 250 cc., 
(c) 500 cc., (d) 721 cc. ? 

6. Express in grams (a) 721 kg., () 62 mg., (c) 245 eg., (d} 84 dg. 

7. Express (a) 40 meters in inches, () 25 kilograms in pounds, 
(c} 54 grams in ounces, (d) 72 grams in grains, (e) 75 liters in quarts 

2. The Thermometer in scientific use is the centigrade. The boil- 
ing point of water on this thermometer is zoo, and the freezing point is 
o (Fig. 80). The equal spaces between these points are called degrees. 
The abbreviation for centigrade is C., and for degrees 
is . Thus, the boiling point of water is 100 C. 
Degrees below zero are always designated as minus, 
e.g. 12 C., means 12 degrees below zero. 

The thermometer in popular use is the Fahrenheit. 
On this instrument the boiling point of water is 212 
and the freezing point is 32 above zero (Fig. 80). 

To change Fahrenheit degrees into the equivalent 
centigrade degrees, subtract 32 and multiply the 
remainder by f , or briefly 

C = f(F- 3 2). 

To change centigrade degrees into the equivalent 
Fahrenheit temperature, multiply by f and add 32 to 
the product, or briefly 


FIG. 80. Ther- 

The point 273 C. is called absolute zero. Absolute temperature 
is reckoned from this point. Degrees on the absolute scale are found 
by adding 273 to the readings on the centigrade thermometer. Thus, 
273 absolute is o C., 274 absolute is + 1 C., etc. 


Descriptive Chemistry. 


1. Change into Fahrenheit readings the following centigrade read- 
ings : (a) 60.5, (J) 40, (0 92, (<0 - 5> (') o, (/)ioo, () 860, (A) -40. 

2. Change into centigrade readings the following Fahrenheit read- 
ings : (a) 207, (b) 1 80, (0 o, (W) -30, (*) 212, (7) 100, () -40, 
(X) 270. 

3. Express the following centigrade readings in absolute readings : 
() o, (*) 24, (*) -13,00 - 26 - 

3. Crystallization. Most substances in passing from a liquid or a 
gas into a solid assume a definite shape. This change is called crys- 
tallization, and the substances are said to crystallize or to form crys- 
tals. Crystals are produced by (i) evaporating a solution, (2) cooling 
a melted solid, or (3) cooling a vapor. Thus, salt crystals are formed 
by evaporating a salt solution ; sulphur crystals, by melting and then 
cooling sulphur, and iodine crystals, by heating iodine in a test tube. 
These methods are called, respectively, evaporation, fusion, and sub- 

As a rule each substance has an individual crystal form by which it 
can be distinguished. Although there are thousands of different crys- 
tals, all belong to one of six classes or systems. This classification is 
based upon two assumptions: (i) all crystals contain certain lines 
called axes, and (2) the surfaces or faces are grouped around the axes 
in definite positions. The axes connect angles, edges, or faces, which 
are similarly situated on opposite sides of the crystal. The bounding 
planes or faces are arranged symmetrically around the axes, which also 
determine (by their lengths and relative positions) the positions of the 
bounding planes. For example, the cube has three equal axes at right 
angles to one another, and terminating in the center of each of the six 
bounding surfaces. 

The following is a brief description of the six 'systems of crystal- 
lization : 

FIG. 81. Isometric crystals (cube, octahedron, dodecahedron). 



(i) Isometric. This has three equal axes intersecting at right 
angles. The simplest forms are the cube, octahedron, and dodecahedron 
(Fig. 81). Substances crystallizing in this system are diamond, com- 
mon salt, alum, fluor spar, iron pyrites, and garnet. 

FIG. 82. Tetragonal crystals. 

(2) Tetragonal. This has three axes at right angles ; but one axis 
is shorter or longer than the other two, which are equal. The common 
forms are the prism, pyramid, and their combinations (Fig. 82). Tin 
dioxide and zircon form tetragonal crystals. 

FIG. 83. Orthorhombic crystals. 

(3) Orthorhombic. This has three unequal axes intersecting at 
right angles. Common forms are the prism, pyramid, and their com- 
binations (Fig. 83). Potassium nitrate, barium sulphate, topaz, and 
native sulphur crystallize in this system (see Fig. 49). 

FIG. 84. Hexagonal crystals. 

(4) Hexagonal. This has four axes : three are equal and intersect 
at 60 in the same plane ; the fourth is longer or shorter than the others 


Descriptive Chemistry. 

and is at right angles to their plane. It is a complex system. Common 
forms are the prism, pyramid, rhombohedron, scalenohedron, and their 
combinations (Fig. 84). In this system are found quartz, calcite, beryl, 
corundum, and ice (see Figs. 5, 52, 61). 

(5) Monoclinic. This has three unequal axes : two cut each other 
obliquely, and the third is at right angles to the plane of the other two. 
Common forms are combinations of prisms. It is a complex system, 
but includes many substances, e.g. sulphur deposited by fusion, sodium 
carbonate, borax, gypsum, and ferrous sulphate (Fig. 85). 

FIG. 85. Monoclinic crystal. 

FIG. 86. Triclinic crystals. 

(6) Triclinic. This has three unequal axes, all intersecting at 
oblique angles. Common forms are complex combinations. Copper 
sulphate, potassium dichromate, boric acid, and several minerals form 
triclinic crystals (Fig. 86). 

4. History and Biography. The biographical data and table 
given here will serve as a basis for this interesting branch of chemistry. 
Additional facts can be obtained from the historical books mentioned 
below (under " Reference Books ") . 

Arrhenius, Svante, 1859 . Swedish physicist. Contributor to 

modern theory of solution. 

Avogadro, Amadeo, 1776-1856. Italian chemist and physicist. Pro- 
posed in 1811 his hypothesis equal number of molecules in equal 
volumes of all gases at same temperature and pressure. 

Balard, Antoine Jerome, 1802-1876. French chemist. Discovered 
bromine in 1826. 

Becher, Johann Joachim, 1635-1682. German physician. Dis- 
covered few facts, but collected and explained writings of others. 
Believed in alchemy, but made no search for gold. Laid foundations of 
phlogiston theory. 

Appendix. 443 

Bergman, Torbern, 1735-1784. Swedish chemist. Improved 
methods of chemical analysis. Believed in phlogiston. Studied min- 
erals and organic acids. Contributed much to the industrial develop- 
ment of Sweden. Intimate friend of Scheele. 

Berthollet, Claude Louis, 1748-1822. French chemist. Studied 
composition of ammonia, properties and nature of chlorine, hydrogen 
sulphide, and hydrocyanic acid. Explained chemical changes by 
" affinity." His discussion with Proust led to law of definite proportions. 

Berzelius, Johann Jacob, 1779-1848. Swedish chemist. Deter- 
mined many atomic weights. Introduced use of symbols. Discovered 
selenium, prepared silicon and several rare elements. Investigated law 
of multiple proportions, proposed dualistic theory and an electrochem- 
ical theory, improved experimental methods. Industrious investigator, 
prolific writer. 

Bessemer, Sir Henry, 1813-1898. English metallurgist. Devised, 
in 1856, Bessemer process of making steel. 

Black, Joseph, 1728-1799. Scotch chemist and physicist. Dis- 
covered carbon dioxide. Showed relation of this gas to carbonates of 
alkalies and alkaline earths. Opposed phlogiston theory. Teacher 
and friend of James Watt and Rutherford. 

Boyle, Robert, 1626-1691. English philosopher. Announced law 
of effect of pressure on gases. Studied air and water. Opposed 
to alchemy. Views anticipated present conception of constitution of 
matter. Laid foundation of qualitative analysis. 

Bunsen von, R. W. E., 1811-1899. German chemist. Studied 
blast furnace and developed gas analysis. Invented the burner, pho- 
tometer, and battery bearing his name. With Kirchhoff (about 1860) 
devised the spectroscope, and by it developed spectrum analysis and 
discovered rubidium and caesium ; improved the calorimeter ; studied 
chemical action of light. 

Cannizzaro, Stanislao, 1826 . Italian chemist. Revived Avoga- 

clro's hypothesis in 1858, and thereby led to revision of atomic weights. 

Cavendish, Henry, 1731-1810. English chemist. Discovered hy- 
drogen, determined specific gravity of gases, showed (i) solubility of 
calcium carbonate in water containing carbon dioxide, (2) formation of 
water by burning of hydrogen. Determined composition of the atmos- 
phere and of nitric oxide. Accepted phlogiston theory. He was 
parsimonious, eccentric, shy ; trained mathematician and electrician ; 
" the richest of the wise, and the wisest of the rich." 

444 Descriptive Chemistry. 

Charles, Jacques Alex Cesar, 1746-1822. French physicist. Pro- 
posed law bearing his name. 

Courtois, Bernard, 1777-1838. French chemist. Discovered iodine 
in 1811. 

Dalton, John, 1766-1844. English chemist, physicist, and mathe- 
matician. Devised atomic theory. Discovered law of multiple propor- 
tions. " Dalton was often inaccurate as to facts, deficient in the details 
of chemical manipulations, and did not hold high rank as an experi- 
menter; but he was good at drawing conclusions and at stating 
generalizations, his aim being the establishment of general, underlying 
laws." (Venable.) 

Davy, Sir Humphry, 1778-1829. English chemist. Studied gases, 
demonstrated properties of nitrous oxide, determined composition of 
hydrochloric acid, studied iodine and chlorine, named latter. Isolated 
potassium, sodium, barium, calcium, and strontium by electrolysis, and 
studied action of electricity on water and on many other substances. 
Devised miner's safety lamp. "He was one of the most brilliant 
chemists the world has ever seen and the greatest England has pro- 

Dewar, James, 1842 . English chemist. Pioneer in the lique- 
faction of gases by modern methods. (See Hydrogen.) 

Dulong, Pierre Louis, 1785-1838. French chemist and physicist. 
With Petit announced law of specific heats in 1819. 

Dumas, Jean Baptiste Andre, 1800-1884. French chemist. Deter- 
mined many atomic weights, gravimetric composition of water, compo- 
sition of air. Investigated many organic compounds. Devised a 
method of determining vapor density. Excellent teacher, careful 
editor, and faithful public servant. 

Faraday, Michael, 1791-1869. English chemist and physicist. 
Liquefied chlorine and other gases. Showed quantitative relation be- 
tween electric current and chemical changes, and developed electro chem- 
istry. Was Davy's assistant and successor in the Royal Institution. 
Popular lecturer, keen investigator, and ardent lover of science. 

Gay-Lussac, Joseph Louis, 1778-1850. French chemist and physi- 
cist. Announced law of gas volumes in 1808. Worked on cyanogen, 
iodine, halogen acids, alkaline oxides, isolation of boron. Improved 
methods of analyzing organic compounds. Was pupil of Berthollet. 
" Was a trained chemist, capable of most accurate analytical work, 
and possessing scientific acumen in a very high degree." (Venable.) 

Appendix. 445 

Glauber, Johann Rudolph, 1604-1668. German chemist. Believed 
in alchemy. Discovered sodium sulphate, which even now bears his 
name. Suggested improvements in industrial chemistry. 

Graham, Thomas, 1805-1869. British chemist. Studied diffusion 
of gases, acids of phosphorus, water of crystallization, and dialysis. 
Developed idea of basicity of acids. 

Hofmann von, August Wilhelm, 1818-1892. German chemist. 
Studied organic chemistry exhaustively. Coal-tar industry arose 
largely from his work. Devised unique lecture apparatus, e.g. that for 
the electrolysis of water. Brilliant teacher, prolific investigator. 

Kirchhoff , Gustav Robert, 1 824-1 887. German physicist. With Bun- 
sen, devised spectroscope and founded principles of spectrum analysis. 

Lavoisier, Antoine Laurent, 1743-1794. French chemist. Over- 
threw phlogiston theory, explained combustion, contributed many facts 
to a large number of chemical topics. Devised foundation of chemical 
nomenclature. Interpreted experiments of other chemists. Efficient 
public servant. Regarded by many as the founder of modern chem- 
istry. Accused of appropriating public money and of " putting water 
in the people's tobacco," he was condemned by the infamous Robes- 
pierre, and publicly guillotined. 

Liebig von, Justus, 1803-1873. German chemist. Laid founda- 
tions of agricultural and organic chemistry. Eminent teacher. 

Mendeleeff, Dmitri Ivanovitch, 1834 . Russian chemist. An- 
nounced periodic law in 1868. 

Meyer, Lothar 1830-1895. German chemist. Contributed to estab- 
lishment of periodic law. 

Moissan, Henri, 1852 . French chemist. Isolated fluorine, 

devised and perfected electric furnace, prepared artificial diamonds, 
rare metals, and refractory compounds. 

Ostwald, Wilhelm, 1853 . German chemist. Contributor to 

modern theory of solution. Eminent teacher and prolific writer. 

Petit, Alexis Therese, 1791-1820. French physicist. (See Dulong.) 

Priestley, Joseph, 1733-1804. English chemist and theologian. 
Student of electricity, light, and gases. Discovered oxygen. Devised 
pneumatic trough. His political and religious views were so freely 
expressed that he was obliged to leave England. Came to America in 
1795. Died at Northumberland near Philadelphia, Pennsylvania. 

Proust, Louis Joseph, 1755-1826. French chemist. Defended 
law of definite proportions in a long controversy with Berthollet. 

446 Descriptive Chemistry. 

" One of the good results of this controversy was to bring about a defi- 
nition of compounds and mixtures, and a clear distinction between 
them. In course of it, also, Proust discovered the hydroxides, a class 
of compounds until then confused with the oxides." (Venable.) 

Prout, William, 1785-1850. English physician. Advanced in 1815 
the hypothesis that the atomic weights of all elements are whole 

Ramsay, William, 1852 . English chemist. Discovered 

argon, helium, neon, krypton, and xenon. 

Rutherford, Daniel, 1749-1819. Scotch botanist and physician. 
Discovered nitrogen in 1772. Pupil of Black. 

Scheele, Carl Wilhelm, 1742-1786. Swedish chemist. Discovered 
chlorine, ammonia, manganese, baryta, many acids (organic and inor- 
ganic), and oxygen (independently of Priestley). Isolated and studied 
borax, glycerine, Prussian blue, microcosmic salt. Improved the 
methods of preparing many substances. Was very poor. Friend and 
companion of Bergman. Achieved marvelous results with simple 
appliances. Believed in phlogiston. 

Stahl, George Ernst, 1660-1734. German physician and chemist. 
Revived and extended Becher's ideas of combustion. Introduced the 
name phlogiston. Strongly advocated this theory. Successful teacher 
and writer. 

Stas, Jean Servais, 1813-1891. Belgian chemist. Determined 
accurately many atomic weights. Pupil of Dumas. Overthrew Prout's 

Van Helmont, Jean, 1577-1644. Dutch chemist. Studied gases, 
and discovered carbon dioxide. Had imperfect but introductory views 
on physiological chemistry, indestructibility of matter, and elements. 
Believed in the alkahest or universal solvent. 

Van't Hoff, Jacobus Hendricus, 1852 . Dutch chemist. Con- 
tributor to chemistry of space relations of atoms and to modern theory 
of solution. 

Wohler, Friedrich, 1800-1882. German chemist. Isolated alu- 
minium and beryllium. Worked on boron, silicon, and many organic 
substances. Discovered isomerism. Overthrew barrier between or- 
ganic and inorganic chemistry. Was fellow-worker with Liebig, pupil 
of Berzelius, and influential teacher of many famous chemists. 

Appendix. 447 







Albertus Magnus Roger Bacon 

8th Century 978-1036 



Middle Ages 

Raymond Lulli 

Basil Valentine 



I4th to i6th Cen- 




Van Helmont 










lyth and i8th 



























i8th and igth 



















Graham FRENCH. Dumas 





igth Century. 










5. Atomic Weights. The following table of atomic weights is 
from the Journal of the American Chemical Society, Vol. XXV, No. I 
(January, 1903). 


Descriptive Chemistry. 











Aluminium .... 














Arsenic .... 





Barium .... 





Bismuth .... 









Bromine .... 





Cadmium .... 




Caesium .... 




Calcium .... 





Carbon .... 





Cerium .... 




Chlorine .... 





Chromium .... 



5 1 -? 






Columbium .... 




Copper .... 





Erbium .... 




Fluorine . . . ' 









Gallium . 








Glucinum .... 




Gold . . . . 4 





Helium .... 


m TM 



Hydrogen . . . . 





Indium .... 









Iridium .... 








Krypton .... 





Lanthanum .... 









Lithium .... 





Magnesium .... 





Manganese .... 





Mercury . . . . 









1 Use these values in solving problems. 




























Osmium .... 




Oxygen .... 




Palladium .... 




Phosphorus .... 





Platinum .... 





Potassium .... 





Praseodymium . . . . 




Radium .... 




Rhodium .... 

' Rh 



Rubidium . . . 




Ruthenium .... 




Samarium ... 




Scandium .... 




Selenium .... 














Sodium . 





Strontium ... 




Sulphur .... 





Tantalum .... 




Tellurium .... 




Terbium .... 




Thallium .... 




Thorium .... 




Thulium .... 









Titanium .... 




Tungsten .... 




Uranium .... 




Vanadium .... 







Ytterbium .... 




Yttrium .... 









Zirconium .... 




1 Use these values in solving problems. 

450 Descriptive Chemistry. 

6. Reference Books and Supplementary Reading. The list of 
books given below will serve as the basis of a chemical library. The 
starred (*) titles indicate books intended for the teacher, though many 
parts of these books are not beyond the grasp of pupils. The library 
should contain at least numbers i, 5, 8, 10, 18, 20, 22, 24. Additional 
titles can be found in (i) List of Books in Chemistry, L. E. Knott 
Apparatus Co., Boston, Mass. ; (2) Smith and Hall's Teaching of 
Chemistry and Physics, p. 218; (3) NEWELL'S EXPERIMENTAL CHEM- 

i. Text-Book of Inorganic Chemistry, Newth. Longmans, Green, 
& Co., 682 pp., $1.75. 

*2. General Inorganic Chemistry, Freer. Allyn & Bacon, Boston, 
559 PP- #3- 

*3. Text-Book of Inorganic Chemistry, Holleman. John Wiley & 
Sons, 458 pp., $2.50. 

4. Physical Chemistry for Beginners, Van Deventer. John Wiley 
& Sons, 154 pp., $1.50. 

5. Chemical Theory for Beginners, Dobbin and Walker. The 
Macmillan Co., 236 pp., $ .70. 

*6. Introduction to Physical Chemistry, Walker. The Macmillan 
Co., 332 pp., $3. 

7. The Birth of Chemistry, Rodwell. The Macmillan Co., 135 
pp., $i. 

8. Short History of Chemistry, Venable. D. C. Heath & Co., 172 
pp., $i. 

9. Faraday as a Discoverer, Tyndall. D. Appleton & Co., 171 
pp., $i. 

10. Short History of Natural Science, Buckley. D. Appleton & 

CO., 467 pp., $2. 

11. Heroes of Science Chemists, Muir. Thomas Nelson & Son, 
350 pp., $1.50. 

*I2. Essays in Historical Chemistry, Thorpe. The Macmillan Co., 
582 pp., $4. 

13. Humphry Davy, Thorpe. The Macmillan Co., 240 pp., $1.25. 

14. John Dalton, Roscoe. The Macmillan Co., 216 pp., $1.25. 

15. Michael Faraday, Thompson. The Macmillan Co., 308 pp., 

*i6. Alembic Club Reprints, University of Chicago Press, $.40 each, 
(i) Experiments on Magnesia Alba. (2) Foundations of the Atomic 

Appendix. 451 

Theory. (3) Experiments on Air. (4) Foundations of the Molecular 
Theory. (6) Decomposition of the Fixed Alkalies. (7) (8) Discov- 
ery of Oxygen. (9) Elementary Nature of Chlorine. (13) Early His- 
tory of Chlorine. 

*ij. Organic Chemistry, Remsen. D. C. Heath & Co., 426 pp., 

18. Outlines of Industrial Chemistry, F. H. Thorp. The Macmil- 
lan Co., 528 pp., $3.50. 

*I9- Practical Electro-Chemistry, Blount. The Macmillan Co., 
374 pp., $3.25. 

20. Chemistry in Daily Life, Lassar-Cohn. J. B. Lippincott Co., 
336pp., $1.75. 

21. The Soil, King. The Macmillan Co., 400 pp. 

22. Story of a Piece of Coal, Martin. D. Appleton & Co., 165 pp., 

23. Chemical History of a Candle, Faraday. Harper & Bros., 223 
pp., $1.00. 

24. Minerals and How to Study Them, E. S. Dana. John Wiley & 
Sons, 380 pp., $1.25. 

*25. Teaching of Chemistry and Physics, Smith and Hall. Long- 
mans, Green & Co., 384 pp., $1.50. 

26. Story of Nineteenth-Century Science, Williams. Harper & 
Bros., 475 PP-> $2.50. 

27. Stories of Industry, Vol. I, Chase and Clow. Educational Pub- 
lishing Co., Boston, 172 pp., $.40. 

Scientific American, Munn & Co., New York. $3.00 yearly; single 
copies, 8 cents. 

School Science, Ravenswood, Chicago, Illinois. $2.00 yearly (9 
issues) ; single copies, 25 cents. 

Popular Science Monthly, The Science Press, New York. $3.00 
yearly ; single copies, 25 cents. 




(Numbers in parentheses indicate experiments.) 


INTRODUCTION . . . . . . ..... . 459 

Bunsen Burner ; Heating ; Cutting and Bending Glass Tubing ; 
Filtering ; Constructing and Arranging Apparatus ; Manipula- 
tion ; Smelling and Tasting. 

PHYSICAL AND CHEMICAL CHANGES . . " . ' . . . . 467 

Physical Change (i, 2, 3); Chemical Change (4). 

OXYGEN . . ... . . . ... . . 468 

Preparation (5) ; Properties (6) ; Preparation from Mercuric Oxide 


HYDROGEN . . . -. . . . . . . ;. . 471 

Preparation (8); Properties (9); Burning Hydrogen (10). 

WATER . . . . . . . . . . . 474 

General Distribution (n); Tests for Impurities (12); Distillation 
(13); Solubility of Gases (14); Solubility of Liquids (15); Solu- 
bility of Solids (16); Supersaturation (17); Water of Crystal- 
lization (18); Efflorescence (19); Deliquescence (20); Solution 
and Chemical Action (21) ; Electrolysis (22) ; Water and Chloripe 
(23); Water and Sodium (24). 

THE AIR , 481 

Composition (25); Water Vapor (26) ; Carbon Dioxide (27). 


Properties of Acids (28) ; Properties of Bases (29) ; A Property of 
Salts (30); Nature of Common Substances (31); Neutralization 

Heat and Chemical Action (33, 34) ; Light and Chemical Action 
(35) ; Electricity and Chemical Action (36, 37). 


456 Descriptive Chemistry. 



Preparation (38) ; Properties (39) ; Bleaching Powder (40) ; Prep- 
aration of Hydrochloric Acid (41); Properties of Hydrochloric 
Acid Gas (42) ; Properties of Hydrochloric Acid (43); Tests for 
Hydrochloric Acid and Chlorides (44). 


Preparation of Ammonia (45); Properties of Ammonia Gas (46); 
Properties of Ammonium Hydroxide (47) ; Neutralization of Am- 
monia (48) ; Preparation of Nitric Acid (49) ; Properties of Nitric 
Acid (50); Test for Nitric Acid and Nitrates (51-52); Interaction 
of Sodium Nitrate and Sulphuric Acid (53) ; Nitric Acid and Metals 
(54); Nitric Acid and Copper, and Nitrogen Peroxide (55); Ni- 
trous Oxide (56); Sodium Nitrite (57); Aqua Regia (58). 

CARBON .,...'. . . 498 

Distribution (59) ; Decolorizing Action (60) ; Deodorizing Action 
(61 ) ; Preparation of Carbon Dioxide (62) ; Properties of Carbon 
Dioxide (63) ; Interaction of Calcium Carbonate and Hydro- 
chloric Acid (64); Carbon Dioxide and Combustion (65); Car- 
bonic Acid (66) ; Carbonates (67) ; Detection of Carbonates (68) ; 
Acid Calcium Carbonate (69); Carbon Monoxide (70); Ethylene 
(71); Acetylene (72); Illuminating Gas (73); Combustion of 
Illuminating Gas (74); Bunsen Burner (75); Bunsen Burner 
Flame (76); Candle Flame (77); Kindling Temperature (78); 
Reduction and Oxidation (79). 


Hydrofluoric Acid (80); Bromine (81); Potassium Bromide (82); 
Iodine (83); Tests for Iodine (84, 85); Detection of Starch 
(86); Potassium Iodide (87). 

SULPHUR . . '> . . . . ... . . . 514 

Properties (88) ; Amorphous Sulphur (89) ; Crystallized Sulphur (90) ; 
Combining Power (91); Sulphur and Matches (92); Preparation 
of Hydrogen Sulphide (93); Properties of Hydrogen Sulphide 
Gas (94) ; Sulphides (95) ; Preparation of Sulphur Dioxide (96) ; 
Properties of Sulphur Dioxide Gas (97) ; Properties of Sulphurous 
Acid (98); Sulphuric Acid and Organic Matter (99); Test for 
Sulphuric Acid and Sulphates (100). 

Contents. 457 



Silicic Acid (101); Borax Beads (102); Boric Acid (103). 


Properties of Phosphorus (104); Test for Arsenic (105); Test for 
Antimony (106); Test for Bismuth (107). 


Properties of Sodium (108); Sodium Hydroxide (109); Exercises; 
Properties of Potassium (no); Potassium Hydroxide (in); Potas- 
sium Carbonate (112); Exercises. 


Properties of Copper (113); Tests for Copper (114); Interaction of 
Copper with Metals (115); Exercises; Preparation of Silver (116); 
Properties of Silver (117); Test for Silver (118); Exercises ; Test 
for Gold (119). 


Tests for Calcium ( 1 20) ; Plaster of Paris (121) ; Exercises; Test for 
Strontium (122); Red Fire (123); Tests for Barium (124); Green 
Fire (125); Exercises. 

Properties of * Magnesium (126); Tests for Magnesium (127); 
Exercises ; Properties of Zinc (128); Tests for Zinc (129); Inter- 
action of Zinc and Metals (130); Exercises; Test for Cadmium 
(131); Properties of Mercury (132); Tests for Mercury (133); 
Mercurous and Mercuric Compounds (134); Exercises. 

ALUMINIUM . . . 532 

Properties (135); Action with Acids and Alkalies (136); Aluminium 
Hydroxide (137); Tests (138); Alum (139). 


Properties of Tin (140); Action of Tin with Acids (141); Tests for 
Tin (142); Deposition (143); Properties of Lead (144); Tests 
for Lead (145); Deposition of Lead (146); Oxides of Lead (147); 
Compounds of Lead (148). 


Tests for Chromium (149); Chromates (150); Reduction of Chro- 
mates (151); Chromic Hydroxide (152); Chrome Alum (153); 
Tests for Manganese (154); Potassium Permanganate (155); 

458 Descriptive Chemistry. 



Properties of Iron (156); Ferrous Compounds (157); Ferric Com- 
pounds (158); Reduction of Ferric Compounds (159); Oxidation 
of Ferrous Compounds ( 1 60); Compounds of Iron (161); Exercises; 
Test for Nickel (162); Test for Cobalt (163). 


Composition (164); Alcohol (165); Properties of Alcohol (166); 
Aldehydes (167); Ether (168); Acetic Acid (169); Vinegar 
(170); Test for Acetic Acid and Acetates (171); Acetates (172); 
Organic Acids (173); Ethyl Acetate (174); Soap (175, 176); 
Glycerine (177); Test for Sugar (178); Exercises ; Benzene (179). 

LABORATORY EQUIPMENT . . ... . . . . 549 

Apparatus ; Chemicals ; Solutions. 


1. The Bunsen burner is used as the source of heat in most chem- 
ical laboratories (Fig. 87). It is attached to the gas cock by a piece of 
rubber tubing. When the gas is turned on, the current of gas draws 
air through the holes at the bottom of the tube, and this mixture when 
lighted burns with an almost colorless, /. e. non-luminous, flame. It is a 
hot flame and deposits no soot. The burner is lighted by turning on 
the gas full and holding a lighted match in the gas about 5 centimeters 
(2 inches) above the top of the burner. If the 
flame is not colorless, or nearly so, turn the ring at 
the bottom of the burner until the flame is a faint 
blue. The colorless flame should be used in all 
experiments unless the directions state otherwise, 
and should be from 5 to 10 centimeters (2 to 4 
inches) high. The hottest part of the flame is near 
the top. 

FIG. 87. Bunsen 

2. Heating. The following directions should 
be observed in heating with the Bunsen burner : 

(1) The burner should always be lighted before 

any piece of apparatus is held over it, or before it is placed beneath 
a wire gauze which supports a dish or flask. 

(2) Glass and porcelain apparatus should not be heated when empty 
nor over a bare or free flame even if they contain something unless 
directions so state. Vessels requiring a support should be placed on a 
wire gauze which stands on the ring of an iron stand, and heated grad- 
ually from beneath. Hot vessels should be heated and cooled gradu- 
ally ; if removed from the gauze while hot, they should be placed on a 
block of wood or piece of asbestos board never on a cold surface. 

(3) Many experiments require the heating of test tubes. These 
tubes should be dry on the outside before being heated. The temper- 
ature of a test tube containing a solid should be raised gradually by 
moving it in and out of the flame, or by holding it in the flame and roll- 




ing it slightly between the thumb and forefinger. Special care must be 
taken to distribute the heat evenly. If the test tube contains a liquid, 
as is usually the case, only that part containing the liquid should be 
heated ; the test tube should also be inclined so that the greatest heat 
is not directed upon the thin bottom. 
When the liquid begins to boil, the test 
tube should be removed from the flame 
for an instant or held over it. In some 
experiments test tubes can be held be- 
tween the thumb and forefinger without 
discomfort. If they are too hot to 
handle, a test-tube holder may be used (Fig. 88). 

3. Cutting and bending Glass Tubing. (a) Cut- 
ting. Determine the length needed, lay the tube on the 
desk, and with a forward stroke of a triangular file make 
a short but deep scratch where the tube is to be cut. 
Grasp the tube in both hands, and hold the thumbs 
together behind the scratch. Now push gently with the 
thumbs, pull at the same time with the hands, and the 
tube will break at the desired point. The sharp ends 
should be smoothed by rubbing them with emery paper 
or by rotating them slowly in the Bunsen flame until a 
yellow color is distinctly seen or until the ends become red-hot. 

() Bending. Glass tubes are bent in a flat flame. 
An ordinary illuminating gas flame may be used, 
but the Bunsen flame can be flattened by a wing-top 
attachment (Fig. 89), which slips over the top of the 
burner tube. 
The flattened 
Bunsen flame 

should be slightly yellow and 
about 7 centimeters (2.5 inches) 
wide for ordinary bends. A 
right-angle bend is made as 
follows: Determine the point 
at which the tube is to be bent. 
Grasp the tube in both hands, 
and hold it so that the part to 

FIG. 88. 
Test tube and 

FIG. 89. Wing- 
top attachment for 
Bunsen burner. 

be bent is directly over the 

FIG. 90. Bending a tube into a right 
angle I. 

Introduction. 461 

flame. Slowly rotate it between the thumbs and forefingers, and 
gradually lower it into the position shown in Figure 90. Continue to 
rotate it until the glass feels soft and ready to yield. Then remove 
it from the flame, and slowly 
bend it into a right angle, as 
shown in Figure 91. It is con- 
venient to have at hand a block 
of wood or some other right- 
angled object to assist the eye 

in completing the bend into an 

. ,. , Tr FIG. 91. Bending a tube into a right 

exact right angle. If a Bunsen angle II. 

flame is used, the bent part of 

the tube should be annealed, i.e. cooled slowly. This is done by 
holding it in a yellow flame until it becomes coated with soot. It 
should then be placed on a block of wood, and when cold wiped 
clean. Tubes can be bent into an oblique angle by heating them 
through about twice the space required for a right angle ; a very slight 
bend, however, is often made by holding the tube across the flame and 
heating a short space. Glass tubes which have been correctly bent 
never have flattened curves ; nor are they twisted, i.e. all parts lie in 
the same plane. 

(c) Drawing. Glass tubes can be drawn to a finer bore or into two 
pointed tubes as follows : Heat the glass as in (b) through about 
2.5 centimeters (i inch) of its length, remove from the flame and 
slowly pull it apart a short distance ; let it cool for a few seconds, and 
then pull it quickly to the desired length. 

The operation is well illustrated by making a glass stirring rod. 
Select a piece of rod about 25 centimeters (10 inches) long and .5 

FIG. 92. Stirring rods ready to be cut. 

centimeter (^ inch) in diameter. Heat it in the middle in the 
ordinary not flat Bunsen flame, and when soft draw it out slowly 
into the shape shown in Figure 92. Cut it into two rods by making a 
slight scratch where the dotted line indicates. Round off the rough 
edges by heating them slightly in the flame. 



4. Filtering. A solid may be separated from a liquid by filtering. 
A circular piece of porous paper is folded to fit a glass funnel, and when 
the mixture is poured upon this paper, the solid the residue or precipi- 
tate is retained, while the liquid the filtrate passes through and 
may be caught in a test tube or any other vessel. The filter paper is 
prepared for the funnel by folding it successively into the shapes shown 
in Figures 93 and 94, and then opening the folded paper so that three 
thicknesses are on one side and one on the other (Fig. 95). The 
cone-shaped paper is next placed in the funnel and wet with water, 

FIG. 93. Folded 
filter paper I. 

FIG. 94. Folded 
filter paper II. 

FIG. 95. Folded filter 
paper ready for funnel. 

so that it will stick to the sides of the funnel and filter rapidly. The 
paper should never extend above the edges of the funnel, but its apex 
should always project slightly into the stem. The liquid to be filtered 
should be poured down a glass rod which touches the edge of the test 
tube ; the lower end of the rod should just touch the paper inside the 
funnel, so that the liquid will run down the side and thereby avoid 
bursting the apex of the filter paper. It is also advisable to adjust the 
apparatus so that the end of the stem of the funnel rests against the 
side of the vessel catching the filtrate. A funnel can be supported by 
standing it in a test tube, a bottle, or the ring of an iron stand. 

5. Constructing and arranging Apparatus. The various parts 
of an apparatus should be collected, prepared, and put together be- 
fore starting the experiment in which the apparatus as a whole is 
used. The different parts which are to fit each other should be selected 
and arranged so that all joints are gas-tight, and as a final precaution 
the apparatus should be tested for leaks. All leaks should be stopped 
up before the apparatus is used. The following hints will be helpful : 

FIG. 96. Rubber tube cut at an angle. 

(1) To insert a glass tube into rubber tubing. Cut the rubber tubing 
at an angle, as shown in Figure 96, moisten the smoothed end of the glass 

Introduction. 463 

tube with water, place the end of the glass tube in the angular-shaped 
cavity so that both tubes are at about a right angle, and then slip the 
rubber tube slowly up and over the end of the glass tube. If the glass 
tube is large or the rubber stiff, the rubber tube must be held firmly 
between the thumb and forefinger to keep it from slipping off until it is 
securely adjusted. 

(2) To fit a glass tube to a stopper. Moisten the end with water and 
grasp the tube firmly about 3 centimeters (i inch) from the end; hold 
the stopper between the thumb and 'forefinger of the other hand, and 
work the tube into the hole by a gradual rotary motion. Proceed in 
the same manner if the tube is to be pushed through the stopper. 
Never point the tube toward the palm of the hand which holds the 
stopper. Never grasp a safety tube or any bent tube at the bend when 
inserting it into a stopper it may break and cut the hand severely. 

(3) To bore a hole in a cork. Rubber stoppers are preferable, but 
if corks are used, they can be bored as follows : Select a cork free from 
cracks or channels and use a borer which is one size smaller than the 
desired hole. Hold the cork between the thumb and forefinger, press 
the larger end against a firm but soft board, and slowly push the borer 
by a rotary movement through the cork, taking care to keep the borer 
perpendicular to the cork. If the hole is too small, enlarge it with a 
round file. If corks are used instead of rubber stoppers, the apparatus 
should always be tested before use by blowing into it, stopping of course 
all legitimate outlets. A poor cork often means a failure, to say noth- 
ing of wasted time. 

(4) To make a platinum test wire. Rotate one end of a piece of 
glass rod, about 10 centimeters (4 inches) long, in the flame until it 
softens. At the same time grasp a piece of platinum wire about 7 cen- 
timeters (3 inches) long firmly in the forceps about i centimeter (.5 
inch) from the end, and hold it in the flame. When the rod is soft 
enough, gently push the hot wire into the rod. Cool the rod gradually 

FIG. 97. Platinum test wire. 

by rotating it in the flame. The completed wire is shown in Figure 97. 
If a glass tube is used instead of a rod, it should be drawn out to a 
very small diameter (see 3 (0) before inserting the platinum wire, 
but in other respects the two operations are practically identical. 



6. Manipulation. Ability to use apparatus rapidly, accurately, and 
neatly is acquired only by experience, but the following suggestions will 
facilitate the acquisition of this needful skill : 

(i) Pouring liquids and transferring solids, (a) Liquids can be 

poured from a vessel without 
spilling, by moistening a glass 
rod with the liquid and then 
pouring it down the rod as 
is shown in Figure 98. The 
angle at which the rod is held 
varies with circumstances. 

This is a convenient way to 
FIG. 98. Pouring a liquid down a glass rod. i- -j r i 

pour a liquid from a vessel 

containing a solid without disturbing the solid. () Liquids can often 
be poured from a bottle by holding the bottle as shown in Figure 99. 
Notice that the stopper and bottle are held in the same hand. This is ac- 
complished by holding the 
palm of the hand upward 
and removing the stopper 
by grasping it between the 
fingers before the bottle is 
lifted. All stoppers should 
be removed this way when 
possible, and not laid down, 
because the impurities ad- 
hering to the stopper may 
run down into the bottle 
and contaminate the solu- 
tion. The drop on the lip of the bottle should be touched with the 
stopper before the latter is put into the bottle ; this simple operation 

prevents the drop from running 
down the outside of the bottle 
upon the label or upon the 
shelf. (<:) Solids should never 
be poured directly from a large 
bottle into a test tube, retort, or 
similar vessel. A convenient 
method is as follows : Rotate 

FIG. 99. The way in which a glass stopper 
should be held while a liquid is being poured 
from a bottle. 

FlG. loo. Pouring a solid into a vessel with 
a small opening. 

the bottle slowly so that the 



solid will roll out in small quantities ; catch the solid on a narrow strip 
of paper folded lengthwise, and slide the solid from the paper into the 
desired vessel. The last part of the operation is shown in Figure 100. 

(2) Collecting gases. Gases are usually collected over water by 
means of a pneumatic trough, a common form of which is shown in Figure 
102. The vessel to be filled with gas is first filled with water, covered 
with a piece of filter paper, inverted, and placed mouth downward on the 
shelf of the trough, which is previously filled with water just above the 
shelf. The paper is then removed, and the vessel slipped over the hole 
in the shelf of the trough. Glass plates instead of filter paper may be 
used to cover the bottle. The gas which is evolved in the generator 
passes through the delivery tube, and bubbles up through the water into 
the bottle, forcing the water out of the bottle as it rises. All gases 
insoluble in water are thus collected. Some heavy gases, such as 
hydrochloric acid, chlorine, and sulphur dioxide, are collected by allowing 
the gas to flow downward into an empty bottle, and displace the air in 
the bottle, i.e. by downward displacement. Ammonia and other light 
gases are usually collected by allowing the gas to flow upward into a 
bottle, i.e. by upward displacement. 

(3) Weighing and measuring. These operations are best learned 
by personal direction from the teacher, together with patient application 
of a few general principles. The following hints, however, will be of 
assistance : 

(a) Learn as soon as possible how to use the scales and interpret the 

(b) Always leave the scales and weights in a clean, usable condition. 

(c) Substances should not be weighed 
on the bare scale pan, but on a smooth 
piece of paper creased on the edges or 
along the middle. Take the solid from 

the bottle with a clean spoon or spatula or --- 

pour by rotating the bottle as described in 

6 (c). In many experiments only ap- -" 12- 

proximate quantities are needed. If you 

weigh out too much, do not put it back 

into the bottle, but throw it away or put it 

into a special bottle. 

(d) Liquids are measured in graduated FlG . IOI .^ Meniscus. Correct 
cylinders. The lowest point of the curved reading is along line I. 

466 Experiments. 

surface of the liquid is its correct height (see Fig. 101). The average 
ordinary test tube holds about 30 cubic centimeters, while the large test 
tube so often mentioned in the succeeding experiments holds about 
75 cubic centimeters. Time can be saved by remembering these volumes. 
(>) All measurements in this book are in the metric system (see App. 
i). The common denominations, their abbreviations, and English 
equivalents should be learned. 

7. Smelling and Tasting. Unfamiliar substances should never 
be tasted or smelled except according to directions, and even then 
with the utmost caution. Never inhale a gas vigorously, but waft it 
gently with the hand toward the nose. Taste acids, etc., by touching a 
minute portion to the tip of the tongue, and as soon as the sensation is 
detected, reject the solution at once never swallow it. 



Experiment 1. Physical Change. Materials: Sugar, glass rod. 

Dissolve a little sugar in a test tube one fourth full of water. Dip a 
glass rod into the liquid and taste it. Has the characteristic property 
of the sugar been changed ? Dip the rod into the liquid again, and 
hold it over the flame of the Bunsen burner. 'As the water evaporates, 
a white solid appears. Taste it. What is it ? Have its original proper- 
ties been destroyed ? What kind of a change did they undergo ? 
What kind of a change did the sugar undergo ? What caused the 
change ? 

Experiment 2. Physical Change. Material: Iodine. 

Drop a small crystal of iodine into a dry test tube, and gently heat 
the bottom. As the violet vapor arises, remove the tube from the flame 
and let it cool. Do the crystals which form near the top resemble the 
original crystal ? When gently heated, do they change into the violet 
vapor ? How has the iodine crystal been changed ? What caused the 
change ? Do the original properties reappear after the cause has been 
removed ? What kind of a change has the iodine undergone ? 

Experiment 3. Physical Change. Material: Glass rod. 

Rub a glass rod briskly on a piece of cloth, and hold it near small 
bits of dry paper. Describe what happens. After a moment touch the 
paper again. Is the result the same ? Try again. Are the original 
properties of the glass restored when the cause of its change is re- 
moved ? What kind of a change did the glass undergo ? 

Experiment 4. Chemical Change. Materials : Copper wire, dilute 
nitric acid, iron nail, forceps. 

(a) Examine a piece of copper wire and notice especially its color. 
Grasp one end of the wire with the forceps, and hold the other end in 
1 467 

468 Experiments. 

the flame until a definite change occurs. Then remove it from the 
flame, and examine. Has it been changed ? Do the original properties 
of the copper reappear when the heated wire is cool ? What kind of a 
change has the copper undergone ? Has the change produced another 
substance ? 

(^) Slip another piece of copper wire into a test tube one fourth full 
of dilute nitric acid. Notice any change. Warm the liquid gently, 
and notice any additional change. What are the evidences of chemical 
change ? What caused the change ? What assisted or hastened it ? 
How has the copper been changed ? (Save the test tube and contents 
for (,).) 

(V) Carefully slip an iron nail into the liquid remaining from ($) ; let 
it stand a short time. Then remove and examine the coating. How 
does it compare with the original copper used in (a) ? What kind of 
a change occurred ? What caused it ? 


(1) What are the evidences of chemical changes in this experiment? 

(2) If a known weight of copper had been consumed in (), could it 
have been obtained without loss in (c) ? 

(3) Did the changes in this experiment involve any loss of copper? 

(4) What is the evidence that new substances were produced in (#) 
and () ? 

(5) What physical changes occurred in (#) and (fr) ? 


Experiment^. Preparation of Oxygen. Materials: jjjjgrarns 
potassium chlorate, 1 5 grams manganese dioxide, g^bottles (about 250 
cubic centimeters each), filter paper, thin piece of soft wood, sulphur, 
deflagrating spoon, piece of charcoal fastened to a wire, piece (about 1 5 
centimeters or 6 inches) of wire picture cord unwound at one end. The 
apparatus is shown in Figure 102. A is a large test tube provided with 
a one-hole rubber stopper, to which is fitted a short glass tube, B ; the 
delivery tube. I), is attached to the short glass tube by the rubber 
tube, C. (Directions for constructing and arranging the apparatus may 
be found in the Introduction, 5.) 

Weigh the potassium chlorate on a piece of paper creased lengthwise, 
and slip it into the test tube ; do the same with the manganese dioxide. 
Shake the test tube until the chemicals are thoroughly mixed ; then hold 



the test tube in a horizontal position and roll or shake it until the mix- 
ture is spread along the tube its entire length. Insert the stopper with 
its tubes, and clamp the test tube to the iron stand, as shown in the 

FIG. 102. Apparatus arranged for preparing oxygen. 

figure, taking care not to crush the tube ; the test tube should incline 
toward the trough, to prevent any water from flowing back upon the 
hot glass. 

Fill the pneumatic trough with water until the shelf is just covered. 
Fill the bottles/)/// of water, cover each with a piece of filter paper, in- 
vert them in the trough, and remove the filter paper ; leave two bottles 
on the shelf and three on the bottom. The end of the delivery tube 
should rest on the bottom of the trough, just under the hole in the shelf. 

Heat the whole test tube gently with a flame about 8 centimeters (or 3 
inches) high. When the gas bubbles regularly through the water, slip a 
bottle over the hole. The gas will rise in the bottle and force out the 
water. Move the flame slowly along the test tube, but concentrate the 
heat toward the closed end, and always keep the flame behind any water 
which may be driven out of the mixture. If the gas is evolved too rapidly, 
lessen the heat; if too slowly, increase it ; if not at all, examine the 
stopper and the rubber connecting tube for leaks, and adjust accordingly. 
When the first bottle of gas is full, remove and cover it with a piece of 
wet filter paper, and slip another bottle over the hole. When five bot- 
tles of gas have been collected, remove the end of the delivery tube 
from the water, lest the cold water be drawn up into the hot test tube 
as the gas contracts. 

Perform the next experiment at once. 

470 Experiments. 

Experiment 6. Properties of Oxygen. 

Proceed as follows with the oxygen prepared in the preceding 

(a) Dip a glowin^_stickjDf^wood into one bottle, and observe the 
change. Remove the stick, and repeat as many times as possible. 
Does the gas burn? How does the glowing stick change? What 
property of oxygen does this experiment show? 

(b} Put a small piece of sulphur in the deflagrating spoon, hold 
the spoon in the flame until the blue flame of the burning sulphur can be 
seen, then lower the spoon into a bottle of oxygen. Notice the change 
in the flame. Describe it. Brush a little of the vapor cautiously 
toward the nose. Of what does the odor remind you? (Plunge the 
spoon into water to extinguish the burning sulphur, and covef the 
bottle with* a piece of filter paper.) 

(c) Hold the charcoal in the flame long enough to produce a faint 
glow, then lower IFmto a bottle of oxygen. Describe the result. 

(a) Melt the sulphur in the deflagrating spoon, and dip the unwound 
end of the wire picture cord into the melted sulphur. Lower the end 
coated with burning sulphur into a bottle of oxygen. The iron wire 
should burn brilliantly. Describe the change. Sometimes the sub- 
stance produced by the change coats the inside of the bottle Describe 
it, if it is visible. 

(tf) With the remaining bottle, repeat any of the above experiments. 


(1) Write a brief account of the above experiments in your note 
book, answering all questions and directions. 

(2) Sketch the apparatus used to prepare oxygen. 

(3) Summarize the properties of oxygen. 

(4) What is its most characteristic property? 

(NOTE. The test tube used in Experiment 5 may be cleaned with 
warm water.) 

Experiment 7. Preparation of Oxygen from Mercuric Oxide. 

Materials : Mercuric oxide, stick of wood. 

Put a little mercuric oxide on the end of a narrow piece of paper 
creased lengthwise, and slip the powder into a test tube. The pow- 
der should nearly fill the round end of the test tube. Hold the test 
tube in a horizontal position, shake it to spread the powder into a thin 



layer, attach the test-tube holder, and heat the test tube (still horizontal) 
in the upper part of the Bunsen flame. Do not heat one place, but 
move the tube back and forth. As soon as a definite change is noticed 
inside the tube, insert a glowing stick of wood. Observe and describe 
the change. If there is no change, heat strongly, and test again. 
.What gas is liberated? Observe the deposit inside the tube. What is 
it? If its nature is doubtful, let the tube cool, and examine again. 


(1) Describe briefly the whole experiment. 

(2) What historical interest has this experiment? 

(NOTE. If the test tube has been partially melted, save it for a sub- 
sequent experiment.) 


Experiment 8. Preparation of Hydrogen. Materials : Granu- 
lated zinc, dilute sulphuric acid, pneumatic trough, four bottles, filter 
paper, taper, matches. The apparatus is shown in Figure 103. A is a 
large test tube provided with a two-hole stopper, through which passes 
the safety tube, B, and the right-angle bend, C\ the long (15 cm. or 6 
in.) delivery tube, E, is attached to the bent tube by the rubber tube, D. 
Precaution. Keep all flames away from the hydrogen generator. 
Fill the test tube half full of granulated zinc as follows : 
Crease a piece of paper lengthwise, pour the zinc from the 
bottle upon the paper, incline the test tube, and slip the zinc 
.into it from the paper do not drop it in. Insert the stopper 
with its tubes ; if the end of the safety tube does not go in 
easily, hold the test tube in a horizontal position and shake 
the zinc about, and at the same time push 
the stopper gently but firmly into place. 
Clamp the apparatus into the position shown 
in the figure or stand it -in a test-tube rack. 
Fill the pneumatic trough 
with water as before, and ad- 
just the apparatus so that the 
end of the delivery tube rests 

on the bottom of the trough 
FIG. 103.- Apparatus for preparing hydrogen. under ^ ho]e ^ ^ ^^ 

Fill the bottles with water and invert them in the trough, as in 
Experiment 5. 


472 Experiments. 

Pour enough dilute sulphuric acid through the safety tube to fill the 
test tube about half full, taking care to leave a little acid in the lower 
bend of the safety tube. This precaution prevents the gas from escap- 
ing from the back of the apparatus ; if at any time the gas should flow 
backward, pour a little acid into the bend ; if the acid does not flow 
down the safety tube, loosen the stopper for an instant. As soon as 
the interaction of the zinc and sulphuric acid produces hydrogen, the 
gas will bubble freely through the water in the trough. Slip a bottle 
over the hole, and collect and remove the bottle of gas as in Experi- 
ment 5, taking care to cover the bottle firmly with a piece of wet filter 
paper. If the evolution of gas slackens or ceases, add a little more acid 
through the safety tube. Collect four bottles of hydrogen, and proceed 
at once with the next experiment. 

Experiment 9. Properties of Hydrogen. 

Study as follows the hydrogen gas prepared above : 

(a} Uncover a bottle for an instant to let a little air in, and then 
drop a lighted match into the bottle. Describe the result. 

() Remove the paper from a bottle of hydrogen, and allow it to 
remain uncovered for three minutes by the clock. Then show the 
presence or absence of hydrogen by dropping a lighted match into the 
bottle. Describe the result. What property of hydrogen is shown by 
this experiment? 

(c} Verify your answer to the last question, thus : Hold a bottle of 
air over a covered bottle of hydrogen, remove the paper, and bring the 
mouths of the bottles close together. (See Fig. i.) Hold them there 
for a minute or two, then stand the bottles on the desk and cover them 
with wet filter paper. Drop a lighted match into each bottle. What 
has become of the hydrogen? How does (c) verify ()? 

(d} Invert a covered bottle of hydrogen, remove the paper, and 
quickly thrust a lighted taper up into the bottle. Withdraw the taper 
and then insert it again. Does the hydrogen burn? If so, where? 
Does the taper burn when in the bottle? When out of the bottle ? 
Feel of the neck of the bottle ; describe and explain. What three 
properties of hydrogen are shown by this experiment ? 

Experiment 10. Burning Hydrogen. (Teacher's Experi- 
ment.) Materials: Apparatus shown in Figure 2, which consists of a 

Hydrogen. 473 

500 cubic centimeter flask fitted with a two-hole rubber stopper, safety 
tube, and double right-angle bend ; the last is attached to a U-tube, 
which is also connected to a delivery tube provided with a short piece 
of capillary glass tubing; calcium chloride, small bottle, platinum wire, 
cotton, granulated zinc, dilute sulphuric acid. 

Fill the U-tube two thirds full of calcium chloride, put a wad of 
cotton beneath the stopper of each arm, and connect the U-tube with 
the generator and the delivery tube. 

Stand the apparatus on the table, examine all joints to be sure they 
are tight, extinguish all flames in the vicinity, and proceed exactly 
according to the following directions : 

Pour slowly but continuously through the safety tube enough (about 
50 cubic centimeters) dilute sulphuric acid upon at least 25 grams of 
granulated zinc to produce a steady current of hydrogen gas for about 
five minutes. It is advisable to use considerable zinc and a moderate 
amount of acid. Acid must not be added after the evolution of gas 
begins, unless, of course, the experiment is begun anew. Let the gas 
bubble through the acid for at least two minutes by actual observation, 
then attach the capillary tube by the rubber connector to the end of 
the delivery tube, leaving a short space between the ends of the two 
glass tubes so that the rubber tube may be compressed suddenly, if 
necessary. Let the gas run for another full minute. This latter pre- 
caution is to drive all air out of the capillary tube. Light the hydrogen, 
and observe at once the nature of the flame, its color, heat (by holding 
a match or platinum wire over it), and any other striking property. 
Then hold a small dry bottle over the flame in such a position that the 
flame is just inside the bottle. When conclusive evidence of the prod- 
uct of burning hydrogen is seen inside the bottle, remove the bottle, 
and extinguish the flame at once by pinching the rubber connector. 
Remove the generator to the hood, and if the evolution of hydrogen is 
still brisk, dilute the acid by pouring water through the safety tube. 
Examine the inside of the bottle. What is the deposit ? Explain its 


(1) What does this experiment suggest about the composition of 
water ? 

(2) Does this experiment illustrate oxidation? Why? Synthesis? 

(3) Describe the whole experiment, and sketch the apparatus. 

474 Experiments. 


Experiment 11. General Distribution of Water. Materials: 
Wood, meat, potato. 

Heat successively in dry test tubes a small piece of wood, of meat, 
or of potato (or any other fresh vegetable) . Hold the open end of the 
test tube lower than the other end. Is there conclusive evidence of 
water? Since most animal and vegetable substances act similarly, what 
general conclusion can be drawn from this experiment ? 

Experiment 12. Simple Tests for Impurities in "Water. 

Materials: Distilled water, water containing dirt, a sulphate, a chloride, 
and a lime compound ; nitric acid, ammonium hydroxide, acetic acid, 
sulphuric acid (concentrated), solutions of potassium permanganate, 
silver nitrate, barium chloride, ammonium oxalate ; and limewater. 

(a) Organic Matter. Fill a clean test tube half full of distilled water, 
and another with water containing a little dirt or a bit of paper. Add 
to each test tube a drop or two of concentrated sulphuric acid and suffi- 
cient potassium permanganate solution (made from distilled water) to 
color each liquid a light purple, as nearly alike as possible. Label one 
tube, and then heat gently nearly to the boiling point the tube contain- 
ing the impure water. As soon as a definite change is seen, heat the 
other cautiously, as too sudden heat may cause the liquid to "bump out." 
Organic matter decolorizes potassium permanganate solution. Which 
tube shows the more organic matter? 

(b) Chlorides. To a test tube half full of distilled water add a few 
drops of nitric acid, and then a few drops of silver nitrate solution. Do 
the same with water known to contain a chloride in solution. What is 
the difference between the results ? The cloudiness, or solid, is due to 
silver chloride, which is always formed when silver nitrate is added to 
hydrochloric acid or a chloride in solution (chlorides being closely related 
to hydrochloric acid). Silver chloride is soluble in ammonium hydroxide. 
Try it. This is the usual test for chlorides (and conversely for soluble 
silver compounds), and will hereafter be used without further description. 

(c} Sulphates. To a test tube half full of distilled water add a few 
drops of sulphuric acid and a few drops of barium chloride solution. 
The white precipitate is barium sulphate. It is insoluble in all common 
liquids, and is always formed when barium chloride is added to sulphu- 
ric acid or a sulphate in solution (sulphates being closely related to sul- 
phuric acid). Test the impure water for sulphates. 

Water. 475 

(y ) Lime Compounds. Add a few drops of a fresh solution of ammo- 
nium oxalate to a test tube half full of clear limewater. Limewater is 
a solution of calcium hydroxide, and the white precipitate formed is 
calcium oxalate, which is soluble in hydrochloric acid but not in acetic 
acid. Try it. This is the test for calcium compounds, often called 
"lime" compounds, because lime, which is calcium oxide, is so well 
known. Apply this test to distilled water and to water known to con- 
tain calcium compounds, and compare the two results. 

(e) Summarize briefly the whole experiment. 

(NOTE. If time permits, this experiment should be applied by the 
class to water whose impurities are unknown.) 

Experiment 13. Distillation. (Teacher's Experiment.) Ma- 
terials: Condenser, etc., shown in Figure 6, potassium permanganate, 
impure water, and solutions used in Experiment 12. 

Fill the flask, C, half full of water known to contain the impurities 
mentioned in Experiment 12, add a few crystals (3 or 4) of potassium 
permanganate, and connect with the condenser as shown in Figure 6. 
Attach the inlet tube to the faucet, fill the condenser slowly, and regu- 
late the current so that a small stream flows continuously from the 
outlet tube into the sink or waste pipe. Heat the liquid in C gradually, 
and when it boils, regulate the heat so that the boiling is not too vio- 
lent. As the distillate collects in the receiver, Z?, test separate portions 
for organic matter, chlorides, sulphates, and calcium compounds. 


(1) Is organic matter found ? 

(2) Is mineral matter found ? 

(3) If the distilling liquid had contained a volatile substance, like 
ammonia or alcohol, would the distillate contain such a substance ? 

Experiment 14. Solubility of Gases. 

(a} Warm a little faucet water in a test tube. Is there immediate 
evidence of a previously dissolved gas ? Is there evidence of much 
gas ? What effect has increased heat ? 

(6) Warm slightly a few cubic centimeters of ammonium hydroxide 
in a test tube. Do the results resemble those in (a) ? As soon as the 
final result is obtained, pour the remaining liquid down the sink and 
flush well with water. 



(V) Repeat (), using a little concentrated hydrochloric acid. Do 
the results resemble those of (a) and (b) ? 

(1) How does increased temperature affect the solubility of gases ? 

(2) What gases dissolve freely in water ? 

Experiment 15. Solubility of Liquids. Materials: Alcohol, 
kerosene, glycerine, carbon disulphide. 

(a) To a test tube half full of water add a little alcohol and shake. 
Is there evidence of solution ? Add a little more and shake. Add a 
third portion. Is there still evidence of solution ? Draw a conclusion 
as to the solubility of alcohol in water. 

() Repeat (), using successively kerosene, glycerine, and carbon 
disulphide. Observe the results and conclude accordingly. 

(c) Summarize the results in a table. 

Experiment 16. Solubility of Solids. Materials: About 20 
grams of powdered copper sulphate, 6 grams of powdered potassium 
chlorate, i gram of calcium sulphate. 

(a) Label three test tubes I, II, III. Fill each about one third full. 
To I add i gram of powdered copper sulphate, to II add i gram of 
powdered potassium chlorate, to III add i gram of calcium sulphate. 
Shake each test tube, and then allow them to stand undisturbed for a 
few minutes. Is there evidence of solubility in each case? Is there 
evidence of a varying degree of solubility? If III is doubtful, carefully 
transfer a portion of the clear liquid to an evaporating dish by pouring 
it down a glass rod (see Introduction, 6 (i )(#)), and evaporate to dry- 
ness. Is there now conclusive evidence of solution? Draw a general 
conclusion from this experiment. Save solutions I and II for (). 

Tabulate the results of (d) as follows, using the customary terms to 
express the degree of solubility : 





i . Copper sulphate 
2. Potassium chlorate 

Water at tempera- 
ture of labora- 



3. Calcium sulphate 



Water. 477 

() Heat I and add gradually 4 more grams of powdered copper 
sulphate. Does it all dissolve? Heat II and add 4 more grams of 
powdered potassium chlorate. Does it all, or most all, dissolve? What 
general effect has increased heat on the solubility of solids? What is 
the difference between this general result and that in Experiment 14? 
Save the solutions for (c) . 

(c) Heat I and II nearly to boiling, and as the temperature in- 
creases add the respective solids. Do not boil the liquid away. Is 
there a limit to their solubility? Draw a general conclusion from these 
typical results. 

Experiment 17. Supersaturation. Material: Sodium thio- 

Fill a test tube nearly full of crystallized sodium thiosulphate and 
add a very little water. Warm slowly. As solution occurs, heat 
gradually to boiling. Add sodium thiosulphate until no more will 
dissolve. Pour the solution into a warm, clean, dry test tube and 
let it stand until cool. Then drop in a small crystal of sodium thio- 
sulphate and watch for any simple but definite change. What hap- 
pens? Is the excess of solid large? How does a supersaturated 
solution differ from a saturated one? 

Experiment 18. Water of Crystallization. Materials: Crys- 
tallized sodium carbonate, gypsum, copper sulphate, evaporating dish, 
gauze-covered ring (or tripod). 

(a) Heat a few small crystals of sodium carbonate in a dry test tube, 
inclining the test tube so that the open end is the lower. What is the 
evidence that they contained water of crystallization? If there is any 
marked change in the appearance of the crystals, describe and explain it. 

(b) Repeat, using a crystal of gypsum. Answer the question asked 
in (a). 

(c} Heat two or three small crystals of copper sulphate in an evapo- 
rating dish which stands on a gauze-covered ring. As the action pro- 
ceeds, hold a dry funnel or glass plate over the dish. Is there conclusive 
evidence of escaping water of crystallization ? Do the crystals change 
in color? In shape? Can the form of the crystals be changed by 
gently touching the mass with a glass rod? Continue to heat until the 
resulting mass is a bluish gray. Let the dish cool. Meanwhile heat a 
test tube one half full of water. When the dish has cooled somewhat, 

478 Experiments. 

pour the hot water slowly into the dish upon the copper sulphate. Ex- 
plain the change in color, if any. If there are any lumps, crush them 
with a glass rod. Let the clear solution evaporate for several hours. 
Are crystals deposited? If not, heat' a few minutes, and cool again. 
If so, why ? Have they water of crystallization, and, if so, where did 
they get it? 

Experiment 19. Efflorescence. 

Put a fresh crystal of sodium carbonate and of sodium sulphate on 
a piece of filter paper, and leave them exposed to the air for an hour or 
more. Describe any marked change. What does this change show 
about the air ? About the crystal ? 

Experiment 20. Deliquescence. 

Put on a glass plate or block of wood a small piece of granulated 
calcium chloride and of sodium hydroxide. Leave them exposed to 
the air for an hour or more. Describe any marked change which takes 
place. Compare the action with that of Experiment 19. 

Experiment 21. Solution and Chemical Action. Materials: 
Powdered tartaric acid, sodium bicarbonate, lead nitrate, potassium di- 
chromate, mortar, dish of water. 

(#) Mix in a dry mortar small but equal amounts of powdered tar- 
taric acid and sodium bicarbonate. Is there any decided evidence of 
chemical action ? Pour the mixture into a dish of water. Is there con- 
clusive evidence of chemical action ? 

(b) Repeat, using powdered lead nitrate and powdered potassium 

Describe the results in (a) and (b). How does solution influence 
chemical action ? Why are so many solutions used in the laboratory ? 

Experiment 22. Electrolysis of Water. (Teacher's Experi- 
ment.) Materials : Hofmann apparatus, sulphuric acid, taper, matches, 
short piece of capillary glass tubing. 

Fill the Hofmann apparatus, Figure 10, with water containing 10 per 
cent of sulphuric acid, so that the water in the reservoir tube stands a 
short distance above the gas tubes after the stopcock in each has been 
closed. Connect the platinum terminal wires with a battery of at least 
two cells. As the action proceeds, small bubbles of gas rise and collect 

Water. 479 

at the top of each tube. Allow the current to operate until the smaller 
volume of gas is from 8 to 10 centimeters in height. Measure the 
height of each gas column. Assuming that the tubes have the same 
diameter, the volumes are in approximately the same ratio as their heights. 
How do the volumes compare ? 

Test the gases as follows : (#) Hold a glowing taper over the tube 
containing the smaller quantity of gas, cautiously open the stopcock to 
allow the water (or air) to run out of the glass tip, and then let out a 
little gas upon the glowing taper. What is the gas ? Repeat until the 
gas is exhausted. Care must be taken not to lose the gas. It is ad- 
visable to have at hand several partially burned tapers or thin splints, 
in case any escaping water extinguishes the first one. (t>) Open the 
other stopcock long enough to force out the water in the glass tip ; 
close the stopcock, and, by means of a short rubber tube, attach the 
capillary tube close to the end of the glass tip. Open the stopcock 
again, let out the gas slowly, and hold at the same time a lighted match 
at the end of the tip, then immediately thrust a taper into the small 
and almost colorless flame. What is the gas ? Repeat until the gas is 


(1) Describe the whole experiment. 

(2) Draw a general conclusion from this experiment. 

(3) What does this experiment show about the composition of 
water ? 

(4) Sketch the apparatus. 

Experiment 23. Interaction of Water and Chlorine. (Teach- 
er's Experiment.) Materials: Glass\ube I meter long and about 2 
centimeters in diameter, cork for one end, evaporating dish, chlorine 

Construct a chlorine generator, as described in Experiment 38, and 
prepare about 250 cubic centimeters of chlorine water by causing the 
gas to bubble through a bottle of water until the water smells strongly 
of the gas. Close one end of the tube with a cork. The cork must 
fit air tight, and as a precaution should be smeared (after insertion) 
with vaseline or coated with paraffin. Fill the tube full of chlorine 
water, cover the open end with the thumb or finger, invert the tube, and 
immerse the open end in the evaporating dish, which should be nearly 

480 Experiments. 

full of chlorine water. Clamp the tube in an upright position, and stand 
the whole apparatus where it will receive the direct sunlight for at least 
six hours. Bubbles of gas will soon appear, rise, and collect at the 
top. When sufficient gas for a test has collected, unclamp the tube, 
cover the open end with the thumb or finger, invert the tube, and put a 
glowing taper into the gas. Repeat as long as any of the gas remains. 


(1) What gas is produced by the interaction of chlorine and water ? 

(2) Describe this experiment. 

(3) What does it show about the composition of water ? 

(4) Sketch the apparatus. 

Experiment 24. Interaction of Water and Sodium. Mate- 
rials : Sodium, pneumatic trough filled with water as usual, tea lead, for- 
ceps, red litmus paper. 

Precaution. Sodium, shotdd be handled cautiously and used strictly 
according to directions. Small fragments must not be left about nor 
thrown into the refuse jar, but into a large vessel of water especially pro- 
vided for that purpose. 

(a) If the sodium is brown, scrape off the coating. Cut off a piece of 
sodium not larger than a small pea, and drop it upon the water in the 
trough. Stand far enough away so that you can just see the action. 
Wait until you are sure the action has stopped, and then describe all you 
have seen. 

(b) The action in (a) may be further studied as follows : Fill a test 
tube with water, invert it, and clamp it in the trough so that the mouth is 
over the hole in the shelf of the trough. Wrap a small piece of sodium 
loosely in a piece of tea lead about 5 centimeters (2 inches) square, make 
two or three small holes in the tea lead, and then thrust it under the 
shelf of the trough with the forceps. A gas will rise into the test tube. 
Proceed similarly with additional small pieces of sodium and dry tea 
lead until the test tube is nearly full of gas ; then unclamp and remove, 
still keeping the tube inverted. Hold a lighted match, for an instant, at 
the mouth of the tube. Observe the result, watching especially the 
mouth of the tube. What is the gas? Why? Remembering that 
sodium is an element, where must the gas have come from? If there is 
any doubt about the nature of the gas, collect more, and subject it to 
those tests which will prove its nature. 

The Air. 481 

(V) Put a piece of filter paper on the water in the trough, and before 
it sinks drop a small piece of sodium upon it. Stand back and observe 
the result. Wait for the slight explosion which usually occurs soon 
after the action stops. Describe all you have seen. What burned? 
What caused it to burn? To what is the vivid color probably due? 
(In answering these questions, utilize your knowledge (i) of the prop- 
erties of the gases previously studied, and (2) of the usual accompani- 
ment of chemical action, suggested here by the melting of the sodium.) 

(d} Test the water in the trough with red litmus paper. Push the 
paper to the bottom or to the place where it is certain that chemical 
action between water and sodium has taken place. Test until the red 
litmus paper has undergone a decided change in color. Describe this 
final result. With another piece of red litmus paper test a solution of 
sodium hydroxide. Is the result similar? Dip a glass rod or the plati- 
num test wire (see Int. 5 (4)) into this solution and hold it in the 
Bunsen flame. Describe the result. Is the color of this flame and that 
noticed in (<:) the same? Are the dissolved substances identical? 

(e) W T hat does the whole experiment show about the composition of 
water ? 


Experiment 25. Composition of the Air. Materials: Solu- 
tions of pyrogallic acid and potassium hydroxide, 1 pneumatic trough half 
filled with water at the temperature of the room, 500 and 25 cubic centi- 
meter graduated cylinders. The apparatus consists of an Erlenmeyer 
flask (250 cubic centimeters) provided with a one-hole rubber stopper 
into (but not through) which passes a short glass tube ; to the outer end 
of this tube, which projects 2.5 centimeters (i inch) above the stopper, 
a rubber tube (5 centimeters or 2 inches long) is tightly fastened ; a 
Hofmann screw is attached to the rubber tube close to the end of the 
glass tube. 

(a) The volume of the flask is found thus : Fill the flask completely 
with water from the pneumatic trough. Loosen the screw and push 
the stopper into the flask as far as it will go. Wipe the flask dry and 
carefully remove the stopper. Pour most of the water from the flask 
into the 500 cubic centimeter graduate, and read the volume : the last 
portions of the water in the flask should be poured into the 25 cubic 

1 The pyrogallic acid is a 10 per cent solution, and the potassium hydroxide 
50 per cent. 

482 Experiments. 

centimeter graduate, so that the volume can be read accurately. (See 
Fig. 101). Record the total volume of the flask as shown in {d}. 

(b) Measure exactly 10 cubic centimeters of pyrogallic acid in the 
small graduate (see Int. 6 (3) (d)), and pour it into the flask. Add 
20 cubic centimeters of potassium hydroxide solution, and insert the 
rubber stopper quickly and firmly. Tighten the screw. Shake the 
flask vigorously for a minute. Then invert it and watch the surface of 
the liquid for bubbles. If any appear, the apparatus leaks. Find 
the leak, if any, start the experiment again from (), taking care to 
remedy the defect before the flask is shaken. If no bubbles appear, 
continue to shake at intervals from fifteen to twenty minutes. During 
this operation the oxygen is absorbed by the solution. 

(c) Place the flask on its side in the water of the pneumatic trough, 
and open the screw, taking care (i) not to let any of the solution run 
out, (2) nor to let too much water run in, and (3) to keep the end of 
the rubber tube constantly below the surface. After the water has 
stopped running in, remove the flask from the trough. Open the flask, 
put a glowing stick into the gas, and observe the result. The gas is 
nitrogen. Measure carefully the volume of the final liquid in the flask. 

(d) Record and calculate as follows : 

(a) Volume of original solution = 30 cc. 

(b) Capacity of flask = cc. 

(c) Volume of air taken (b a) = , 

(d) Final volume of liquid = 

(e) Volume of water which entered (d a) 
(f ) Per cent of water which entered (e -s- c) 

But the per cent of entering water equals the per cent of gas ab- 
sorbed, hence 

(g) Per cent of oxygen 

(h) Per cent of nitrogen (100 g) = 

Experiment 26. Air contains Water Vapor. 

Prove by an experiment that air contains water vapor. 

Experiment 27. Air contains Carbon Dioxide. 

(a) Expose a small bottle of limewater to the air. After a short 
time, examine the surface of the liquid. Describe the change. Ex- 
plain it. 

Acids, Bases, and Salts. 483 

() If a blast lamp (or bicycle pump) is available, replace the lamp 
with a glass tube, and force air through a bottle half full of limewater, 
until a definite change occurs. Describe it. Explain it. 


Experiment 28. General Properties of Acids. Materials : 
Dilute sulphuric, nitric, and hydrochloric acids, glass rod, litmus paper 
(both colors), zinc. 

Fill separate test tubes one third full of each of the acids. Label the 
tubes in some distinguishing manner. 

(a) Dip a clean glass rod into each acid and cautiously taste it. 
Describe the taste by a single word. 

(b) Dip a clean glass rod into each acid and put a drop on both kinds 
of litmus paper. The striking change is characteristic of acids ; draw a 
general conclusion from it. 

() Slip a small piece of zinc into each test tube successively. If no 
chemical action results, warm gently. Test the most obvious product 
by holding a lighted match inside of each tube. What gas comes from 
the hydrochloric and sulphuric acids? 

(d} Summarize the general results of this experiment. 

Experiment 29. General Properties of Bases. Materials: 
Litmus paper (both colors), glass rod, sodium hydroxide and potassium 
hydroxide solutions, and ammonium hydroxide. 

(a) Rub a little of each liquid between the fingers, and describe the 
feeling. Cautiously taste each liquid by touching to the tip of the tongue 
a rod moistened in each, and describe the result. 

(b) Test each solution with litmus paper. Describe the result. 

(c) Summarize the general results of this experiment. 

(d) Compare acids and bases as to taste and to reaction with litmus. 

Experiment 30. A Property of Many Salts and All Neutral 
Substances. Materials: Litmus paper (both colors), glass rod, dilute 
solutions of sodium chloride, potassium nitrate, potassium sulphate, and 
barium chloride. 

Test each solution with litmus paper. Describe the result. Com- 
pare with the litmus reaction of acids and bases. 
Draw a general conclusion from this experiment. 



Experiment 31. The Nature of Common Substances. 

Determine by the litmus test the nature of lemon juice, vinegar, 
sweet and sour milk, washing soda, borax, wood ashes, faucet water, 
baking soda, sugar, cream of tartar, the juice of any ripe fruit and any 
green fruit. 

Make a solution of each of the solids before testing. Tabulate the 
results as follows : 





Experiment 32. Neutralization. Materials: Sodium hydrox- 
''ide (solid), hydrochloric acid, nitric acid, silver nitrate solution, blue 
litmus paper, glass rod, evaporating dish, gauze-covered ring. 

Dissolve a small piece of sodium hydroxide in an evaporating dish 
half full of water. Slowly add dilute hydrochloric acid, until a drop 
taken from the dish upon a glass rod reddens blue litmus paper. Then 
evaporate to dryness by heating over a piece of wire gauze supported 
by a ring. Since the residue mechanically holds traces of the excess 
of hydrochloric acid added, it is necessary to remove this acid before 
applying any test. Heat the dish until all the yellow color disappears, 
then moisten the residue carefully with a few drops of warm water and 
heat again to remove the last traces of acid. This precaution is essen- 
tial to the success of the experiment. 

Test a portion of the residue with litmus paper to find whether it has 
acid, alkaline, or neutral properties. Taste a little. Test (a} a solu- 
tion of the residue for a chloride, and (b} a portion of the solid residue 
for sodium. (See Exps. 12 (t>) and 24.) Draw a definite conclusion 
from the total evidence. 

Heat, Light, Electricity, and Chemical Action. 485 


Experiment 33. Heat and Chemical Action. Materials : 
Lime, evaporating dish, match. 

Put a small piece of lime in an evaporating dish, and sprinkle a little 
water over it. Watch for a change. If no marked change soon occurs, 
add a little more water. Describe the change. Touch a match to the 
mass. Is there evidence of much heat? What caused the heat? 

Experiment 34. Heat and Chemical Action. Materials: Sul- 
phur, powdered iron, dilute hydrochloric acid. 

Put about 3 grams of sulphur and 3 grams of powdered iron in a 
test tube. Cover the mouth of the test tube with the thumb and 
shake until the two substances are well mixed. Attach the test tube to 
the holder and heat strongly in the flame. As soon as the sulphur 
melts and boils and the contents give evidence of decided chemical 
action, remove the test tube at once from the flame, and watch the 
change. Is there evidence of heat? Of increasing heat? Of much 

When the tube is cool, break the end, and examine the contents. De- 
scribe it. It is a compound called iron sulphide, and is the product of 
the chemical action which was started by heat. But the chemical action 
itself was so vigorous that it increased_the heat. 

The fact that the product differs from the original mixture may be 
shown as follows : Add dilute hydrochloric acid to a part of the product 
and also to a little of the original mixture, testing the gaseous product 
in each case by the odor. Is the odor the same? 

State briefly how heat and chemical action are related, using this 
experiment as an illustration. 

Experiment 35. Light and Chemical Action. Materials : 
Potassium bromide, silver nitrate solution, funnel, filter paper, glass rod. 

Dissolve a crystal of potassium bromide in a test tube one fourth full 
of water, add an equal volume of silver nitrate solution, and shake. The 
precipitate is silver bromide. Describe it. Filter (see Int. 4). Remove 
the filter paper from the funnel, unfold it, and expose the silver bromide 
for a few minutes to the light sunlight, if possible. Describe the 
change. What caused the change? How is this property of silver 
bromide utilized ? 

486 Experiments. 

Experiment 36. Electricity and Chemical Action. (Teacher's 

Repeat Experiment 22. 

(1) Define electrolysis, electrode, electrolyte, ion, anion, cation. 

(2) State briefly the accepted explanation of the electrolysis of water. 

(3) Is hydrogen an anion or cation? At what electrode does it 
collect ? 

(4) Answer the same questions (as in 3) about oxygen. 

Experiment 37. Electricity and Chemical Action. (Teach- 
er's Experiment.) Materials : Starch, potassium iodide, mortar and 
pestle, filter paper, sheet tin (or iron), battery of two or more cells. 

Grind together in a mortar a lump of starch and a crystal of potas- 
sium iodide. Add enough water to make a thin liquid. Dip a strip of 
filter paper into the mixture, and spread the wet paper upon a sheet of 
tin (or iron) . Press the end of the wire attached to the zinc (of the 
battery) upon the tin, and draw the other wire across the sheet of 
paper. The marks are caused by iodine which is liberated from the 
potassium iodide and colors the starch. 


(1) Describe briefly this experiment. 

(2) Iodine is a non-metal. At what electrode is it liberated? Is 
iodine an anion or a cation ? 


{Do not inhale chlorine?) 

Experiment 38. Preparation of Chlorine. Materials: Con- 
centrated hydrochloric acid, 30 grams manganese dioxide, bundle of 
fine brass wire, strip of calico, paper with writing in lead pencil and in 
ink, litmus paper (both colors), taper. The apparatus is shown in Fig- 
ure 104. It is the same as that used to prepare hydrogen ; and there 
are also needed four bottles, a wooden block (about 10 centimeters or 
4 inches square) with a hole in the center, and four glass plates to 
cover the bottles. 

Weigh the manganese dioxide upon a piece of paper creased length- 
wise. Slip it into the test tube, A (see Int. 6 ()). Arrange the appa- 



ratus as shown in the figure. Pour enough concentrated hydrochloric 

acid through the safety tube to cover the man- 
ganese dioxide. Heat gently with a small 

flame, keeping the flame below the level of the 

contents of the test tube. Chlorine is rapidly 

evolved as a greenish gas, and passes into the 

bottle, G, which should be removed when full 

(as seen by the green color) and covered with 

a glass plate ; the bottle may be easily removed 

by holding the block, F, in one hand and 

pulling the bottle, G, aside, bending the whole 

delivery tube at the same time at the rubber 

connection, D. If the evolution of gas 

slackens, add more acid through the safety 

tube. Collect four bottles, and perform the 

next experiment at once. FIG. 104. Apparatus 

arranged for preparing 
Experiment 39. Properties of Chlorine. 

Study as follows the gas prepared above : 

(a) Heat the bundle of brass wire and thrust it into a bottle of chlo- 
rine. Describe the result, especially the evidence of chemical action 
and of new products. 

(b) Into a bottle of dry chlorine put a piece of 
calico, litmus paper (both colors), and paper contain- 
ing writing in black and in red ink. Allow the whole 
to remain undisturbed for a few minutes and then 
describe the change, if any. Add several drops of 
water, and describe the change. Draw a general con- 
clusion from the whole experiment. 

(c} Hold a burning taper in a, bottle of chlorine 
long enough to observe the result. Draw a conclusion. 
Verify it thus : Fold a strip of filter paper (about 10 
centimeters or 4 inches wide) into the shape shown 
FIG. 105. Fluted ln Figure 105; cautiously heat 1 about 10 cubic centi- 
paper. meters of turpentine in a large test tube; saturate 

1 Hold the test tube with the holder. Remember that turpentine ignites easily. 
If the turpentine catches fire, press a damp towel over it. 

488 Experiments. 

the paper with the hot turpentine and drop it into a bottle of chlorine. 
Describe the result. When the action is over, examine the paper, and 
draw a conclusion regarding the action between hot turpentine and 

Wax (in the taper) and turpentine are mainly compounds of hydro- 
gen and carbon. Explain the result in (c) . 


(1) Many metals act like the brass in (#). What general conclu- 
sion can be drawn about the reaction of chlorine and metals? 

(2) What is essential for the bleaching action of chlorine? 

(3) What does (c} show about the attraction between chlorine and 
hydrogen ? 

(4) What class of chemical changes is illustrated by (#) ? What 
classes by (c) ? 

(5) What class of chemical changes is illustrated by the preparation 
of chlorine? 

(6) What three striking properties has chlorine? How can it be 
distinguished from all gases previously studied? 

Experiment 40. Bleaching by Bleaching Powder. Mate- 
rials : Bleaching powder, sulphuric acid, calico. 

Put a little bleaching powder into a test tube and add enough water 
to make a thin paste. Add a few drops of dilute sulphuric acid, and 
then dip a strip of bright-colored calico into the mixture. Remove the 
calico in a few minutes, and wash it with water. Describe the change 
in the calico. 

Experiment 41. Preparation of Hydrochloric Acid. Mate- 
rials : The apparatus used in Experiment 38 ; 20 grams sodium chlo- 
ride, concentrated sulphuric acid, pneumatic trough filled with water as 
usual, stick of wood, litmus paper (blue), ammonium hydroxide. 

(a) Put 8 cubic centimeters of water in a small bottle or evaporating 
dish, cautiously add 12 cubic centimeters of concentrated sulphuric acid, 
and stir until the two are mixed. While this mixture is cooling, weigh 
the salt, slip it into the test tube, and then arrange the apparatus as 
shown in Figure 104. Pour half the cold acid mixture through the 
safety tube, let it settle through the salt, and then add the remaining 
acid. Heat gently with a low flame, as in the preparation of chlorine. 
Hydrochloric acid gas ' is evolved, and passes into the bottle, which 

Chlorine. 489 

should be removed when full, as directed under chlorine. A piece of 
moist blue litmus paper held at the mouth of the bottle will show when 
it is full. Collect these bottles, cover each with a glass plate, and set 
aside until needed. 

(b) As soon as the third bottle of gas has been collected, removed, 
and covered, put in its place a bottle one fourth full of water. Adjust 
its height (if necessary) by wooden blocks so that the end of the 
delivery tube is just above the surface of the water. Continue to heat 
the generator at intervals, and the gas will be absorbed by the water. 
Shake the bottle occasionally. 

Meanwhile study the gas already collected. 

Experiment 42. Properties of Hydrochloric Acid Gas. 

Proceed as follows with the hydrochloric acid gas prepared by 
Experiment 41 : 

(#) Insert a blazing stick of wood into a bottle. Remove as soon as 
the change is noticed. Describe the change. Compare the action with 
the behavior of hydrogen and of oxygen under similar conditions. 

(^) Hold a piece of wet filter paper near the mouth of the same 
bottle. Describe the result. What is the cause? 

(c) Invert a bottle, and stand it upon the shelf of the pneumatic 
trough. Describe any change noticed inside the bottle after a few 
minutes. What property of the gas does the result illustrate ? Verify 
the observation by a simple test applied to the contents of the 

(d) Drop into the remaining bottle of gas a piece of filter paper wet 
with ammonium hydroxide. Describe the result. What name has 
the product? 

(e) State other properties of hydrochloric acid gas which you have 
observed ; e.g. color, odor, density. 

Proceed at once with the next experiment. 

Experiment 43. Properties of Hydrochloric Acid. 

Remove the bottle in which the hydrochloric acid gas is being ab- 
sorbed (see Exp. 41 ()), and study the solution as follows : 

(a) Determine its general properties, e.g. taste (cautiously), action 
with litmus, and with zinc. 

(^) Add to a test tube half full of the hydrochloric acid a few drops 
of nitric acid and of silver nitrate solution. The white, curdy precipitate 



is silver chloride. Filter part of the contents of the test tube, and ex- 
pose the precipitate to the sunlight. Describe the change which soon 
occurs. To the remaining contents of the test tube add ammonium 
hydroxide, and shake. Describe the result. 

Experiment 44. Tests for Hydrochloric Acid or a Chloride. 

(a) What is a simple test for hydrochloric acid gas or for concen- 
trated hydrochloric acid ? 

{b} What is the usual test for hydrochloric acid ? 

(c) Dissolve a little sodium chloride in a test tube half full of water, 
and apply the test designated in (). (Suggestions. See Exps. 12 (b) 
and 43 ().) 


Experiment 45. Preparation of Ammonia. Materials : 1 5 grams 
lime, 15 grams ammonium chloride, 3 bottles, 2 glass plates, pneu- 
matic trough filled as usual, litmus paper, stick of 
wood, filter paper. The apparatus is shown (in 
part) in Figure 106. The large test tube, A, is 
provided with a one-hole rubber stopper to which 
is fitted the right-angle bend, C, connected with a 
short glass tube, B (12 centimeters or 5 inches 
long), by the rubber tube, D. 

(a) Weigh the lime and ammonium chloride 
separately, mix them thoroughly on a piece of 
paper, and slip the mixture into the test tube to 
which a little water has been previously added. 
Add a little water. Quickly insert the stopper 
with its tubes, and clamp the test tube as shown 
in the figure (taking care not to crush the test 
tube) . 

Slip the glass delivery tube, B, into a bottle, 
invert the bottle, and hold it so that the tube is 
in the position shown in the figure. Heat the 
test tube gently with a low flame, beginning near the top of the mix- 
ture and gradually working downward. Ammonia gas will pass up 
into the bottle, which should be removed when full and covered with 
a glass plate. A piece of moist red litmus paper held near the mouth 
will show when the bottle is full. Do not smell at the mouth of the 
bottle. Collect two bottles and set aside until needed. 

FIG. 106. Appara- 
tus for preparing and 
collecting ammonia gas. 

Compounds of Nitrogen. 491 

(b) As soon as the last bottle has been collected, rearrange the appa- 
ratus to absorb the ammonia gas in water, as in the case of hydrochloric 
acid (see Exp. 41 (^)). Replace the short glass tube by the delivery 
tube, E, which should pass through the wooden block, f y into a bottle, 
G, one fourth full of water, so that the end is just above the surface of the 
water (see Fig. 104). Continue to heat the generator at intervals, and 
the gas will be absorbed by the water. Shake the bottle occasionally. 

While the solution is being prepared, study the gas already collected. 

Experiment 46. Properties of Ammonia Gas. 

Proceed as follows with the ammonia gas prepared in Experiment 
45 :- 

(a) Test the gas in one bottle with moist litmus paper and with a 
blazing stick. Describe the result. Compare the action with the 
behavior of hydrogen, oxygen, and hydrochloric acid gas, under similar 

() Invert the same bottle and stand it upon the shelf of the pneu- 
matic trough. Describe any change noticed inside the bottle. What 
property of the gas is revealed ? Is it a marked property ? Test the 
contents of the bottle with litmus paper (both colors). 

(c) Pour a few drops of concentrated hydrochloric acid into an 
empty, warm, dry bottle. Roll the bottle until the inside is well coated. 
Cover it with a glass plate, invert it, and stand it upon a covered bottle 
of ammonia gas. Remove both plates at once, and hold the bottles 
together by grasping them firmly about their necks. Describe the 
action, giving all the evidence of chemical action. What is the white 
product ? 

Experiment 47. Properties of Ammonium Hydroxide. 

Remove the bottle in which the ammonia gas is being absorbed 
(see Exp. 45, ()), and study the resulting ammonium hydroxide as 
follows : , 

(a) Determine the general properties, e.g. taste and odor (cau- 
tiously), feeling, action with litmus. 

(b) Warm a little in a test tube. What gas is evolved? 

(c) Try the effect of ammonium hydroxide on a grease spot. Describe 
the result. 

Experiment 48. Neutralization of Ammonia. Materials: 
Ammonium hydroxide, hydrochloric acid, evaporating dish, sodium 
hydroxide solution, litmus paper, gauze-covered ring. 

492 Experiments. 

Fill an evaporating dish one fourth full of ammonium hydroxide, and 
slowly add dilute hydrochloric acid, stirring constantly, until the 
solution is just neutral or faintly acid. Evaporate to dry ness, very 
slowly, on a gauze-covered ring. Test the residue as follows : 

(#) Is it an acid, alkali, or salt ? 

(^) Warm a little with sodium hydroxide solution. What is formed? 
Draw a conclusion as to the nature of the residue. 

(c} Support the dish on the gauze and warm gently until a decided 
change occurs. Describe the result. What compound do the fumes 

(</) Verify the observations and conclusions by repeating (3) and 
(c) with ammonium chloride from the laboratory bottle. 

(Y) What is the main product of the interaction of ammonium 
hydroxide and hydrochloric acid? 

Experiment 49. Preparation of Nitric Acid. Materials: 
Glass stoppered retort, sand bath pan and sand, bottle, 30 grams sodium 
nitrate, concentrated sulphuric acid, funnel. 

Weigh the sodium nitrate and slip it into the retort ( see Int. 6 (i) 
(c} ). Attach the retort by a clamp to an iron stand so that (i) Us bulb 
rests on the sand bath, supported by a ring, and (2) the end of its neck 
passes into an inclined bottle which rests on the table. The nitric acid 
which is generated in the bulb will pass down the neck and condense, 
partly in the neck and partly in the bottle. The bottle should be partially 
covered with a piece of wet filter paper, especially where the neck of the 
retort enters. It is advisable, though not always necessary, to place a 
block of wood against the bottom of the bottle to keep it in the desired 

Slip a funnel through the tubulure of the retort as far as it will reach, 
and pour the acid through the funnel into the retort. Remove the funnel 
and insert the stopper of the retort tightly. Heat gently. Brown 
fumes will appear in the retort, and nitric acid will pass into the receiver. 
Distil at as low a temperature as possible, as long as any nitric acid runs 
down the neck of the retort. 

Pour the nitric acid into a test tube or small bottle for use in Experi- 
ment 50. 

Allow the contents of the retort to cool, add a little warm water, let 
the whole stand until the contents are loosened, and then, pour into a 
bottle for use in Experiment 53. 

Compounds of Nitrogen. 493 

Experiment 50. Properties of Nitric Acid. Materials : Quill 
toothpick, indigo solution. Add twice its volume of water to the nitric 
acid prepared in Experiment 49, and proceed as follows : 

(a) Boil a piece of a quill toothpick in a portion of this diluted nitric 
acid. How is the quill changed at first? What is the effect of contin- 
ued heating? Pour off the acid, and wash the quill with water. Is the 
color permanent ? 

(b) Add a dozen or more drops of nitric acid to a dilute solution of 
indigo. Describe the change. Will ammonium hydroxide restore 
the original color? Is the change temporary or permanent? What, in 
all probability, is the general character of the change combination or 
decomposition? Draw a general conclusion from (a) and () regard- 
ing the action of nitric acid on organic matter, which is typified by the 
quill and indigo. 


(1) What color has nitric acid? 

(2) Examine a bottle of nitric acid which has been standing in the 
laboratory. What can be said of the stability of nitric acid? 

(3) State other properties of nitric acid you have observed. 

Experiment 51. Test for Nitric Acid and Nitrates. Materials : 
Concentrated nitric and sulphuric acids, ferrous sulphate, sodium nitrate. 

To a test tube one fourth full of water add a little concentrated nitric 
acid and shake. Add an equal volume of concentrated sulphuric acid. 
Shake until the acids are well mixed, then cool by holding the test tube 
in running water. Make a cold, dilute solution of fresh ferrous sulphate 
and pour this solution carefully down the side of test tube upon the 
nitric acid mixture. Where the two solutions meet, a brown or black 
layer will appear, consisting of a compound formed by the interaction 
of the nitric acid and the ferrous sulphate. It is an unstable com- 
pound and will often decompose if the test tube Is shaken. Record 
the observation. 

This test is also used for a nitrate. Try it with a solution of sodium 
nitrate. Record the result. 

Experiment 52. A Special Test for Nitrates. Materials : Char- 
coal, block of wood, potassium nitrate. 

Heat a piece of charcoal in the Bunsen flame, lay it on a block of 
wood or an iron pan, and cautiously sprinkle powdered potassium ni- 

494 Experiments. 

trate upon the hot surface. Stand back when the action begins. Ob- 
serve and describe the action, especially its violence and rapidity, also 
the color of the flame, the effect on the charcoal, and any other charac- 
teristic result. This kind of chemical action is called deflagration. 
What causes it? 

Experiment 53. The Solid Product of the Interaction of So- 
dium Nitrate and Sulphuric Acid. Materials : Residue from Ex- 
periment 49, evaporating dish, glass rod, gauze-covered ring, distilled 
water, barium chloride solution, ferrous sulphate, concentrated sulphuric 

Pour the solid residue obtained in Experiment 49. into an evaporating 
dish, and evaporate to dry ness over a piece of wire gauze in the hood. 
As the mass approaches pasty consistency, lessen the heat to avoid 
spattering. When the mass is dry, heat strongly as long as white, 
choking fumes are evolved. This last operation is done to remove all 
traces of sulphuric acid, and to complete the chemical change. Allow 
the dish to cool gradually, and when cool, dissolve some of the white 
solid in distilled water and test separate portions for a sulphate and ni- 
trate (see Exps. 12 (V) and 51). Which is it? Test another portion 
for sodium (see Exp. 24 (</) ). What is the name of the white sub- 
stance ? 

Draw a general conclusion regarding the chemical action which oc- 
curs in the preparation of nitric acid by the interaction of sulphuric 
acid and sodium nitrate. 

Experiment 54. Interaction of Nitric Acid and Metals. Mate- 
rials : Zinc, copper, tin, iron, concentrated nitric acid. 

Stand four test tubes in the test-tube rack, and slip into each a few 
small pieces of one of the following metals : zinc, copper, tin, and iron. 
Add to each test tube in succession enough concentrated nitric acid to 
cover the metal. Observe the changes, particularly (i) the vigor of the 
action, (2) the nature and properties of the products, especially color 
and solubility, and (3) evidence of the evolution of hydrogen. Tabu- 
late these observations. 

Experiment 55. Interaction of Nitric Acid and Copper, and 
Study of Nitric Oxide and Nitrogen Peroxide. Materials: 10 
grams copper (borings or fine pieces of sheet metal), concentrated 
nitric acid, pneumatic trough filled as usual, three bottles, three glass 

Compounds of Nitrogen. 495 

plates, matches, piece of wire ( 1 5 centimeters or 6 inches long) ; and 
the apparatus used in Experiment 8. 

Arrange the apparatus as in Experiment 8, after putting the copper 
into the test tube (see Fig. 103). Insert the stopper tightly, adjust 
the delivery tube, fill three bottles with water, and invert them in the 
trough. Pour just enough concentrated nitric acid through the safety 
tube into the flask to cover the copper, taking care to seal the bend of 
the safety tube with acid. Dense brown fumes are evolved. If the 
action is too vigorous, add a little water through the safety tube. Col- 
lect three bottles of the gas which bubbles from the delivery tube. 
Cover them with glass plates and stand them aside until needed. Pour 
the blue liquid in the test tube into an evaporating dish, and evapo- 
rate slowly to crystallization (not to dryness) on a gauze-covered ring in 
the hood. The crystals, after being dried between filter paper, should 
be preserved in a well-stoppered bottle. 

While the solution is evaporating, study the gas as follows : 

(a) Observe its general properties while covered. 

(b} Uncover a bottle. Describe the result. Is the brown gas iden- 
tical in color with the one observed in the generator at the beginning 
of the experiment? 

(c) Uncover a bottle, pour in about 25 cubic centimeters of water, 
cover with the hand and shake vigorously, still keeping the bottle 
covered. Why has the brown gas disappeared? Uncover the bottle 
for an instant, then cover and shake again. Is the result the same? 
Repeat, if the result is not definite, or does not agree with previous 

(d} With the third bottle determine whether the two gases will 
burn or support combustion. A convenient flame is a burning match 
fastened to a stiff wire. Plunge it quickly to the bottom at first and 

gradually raise it into the brown gas. 


(1) What is the source of the colorless gas? What is its name? 
What is the name of the brown gas? 

(2) What is the general chemical relation of the two gases to each 
other? To the air ? 

(3) Why is not the brown gas collected in the bottles by displace- 
ment of water? 

(4) Will either gas burn or support combustion? 

496 Experiments. 

(5) Which gas has been observed before? In what experiment? 

(6) What is the general relation of these gases to nitric acid? 
Study the properties of the crystals by determining : 

(a) Solubility in water (cold and hot). 
() Action of heat. 

(c} Action of their solution upon an iron nail. 
(d) Action of their solution when added to ammonium hydroxide. 
(tf) Presence of a nitrate. 

Compare the observed properties with those of copper nitrate ob- 
tained from the laboratory bottle. Are the two substances identical? 

Experiment 56. Preparation and Properties of Nitrous Oxide. 

Materials : Ammonium nitrate, pneumatic trough filled, as usual, with 
warm water, three bottles, three glass plates, sulphur, deflagrating 
spoo'n, stick of wood. The apparatus is shown in Figure 107. The 
parts lettered A, C, D, E have been used before ; B+ F, G, H are 
exactly the same as A, C, D, E respectively. (See page 504.) 

Construct and arrange the apparatus as shown in the figure. Fill 
the large test tube, A, about half full of ammonium nitrate. The large 
test tube, B, remains empty. The end of //rests on the bottom of the 
pneumatic trough as usual. It is desirable, though not absolutely 
necessary, to fill the trough and bottles with warm water. Be sure the 
apparatus is gas-tight. 

Heat A gently with a low flame (5 centimeters or 2 inches). Adjust 
the apparatus if it leaks. The ammonium nitrate melts and appears to 
boil. Regulate the heat so that the evolution of the nitrous oxide will be 
slow. Notice the fumes which collect in A, and the liquid which col- 
lects in B. Prepare three bottles of nitrous oxide, free from air, cover- 
ing each with a glass plate as soon as removed from the trough. When 
the last bottle has been collected and covered, remove the end of the 
delivery tube from the trough. 

Test the gas as follows : 

(a) Allow a bottle to remain uncovered for a few seconds. How 
does nitrous oxide differ from nitric oxide ? 

(b) Thrust a glowing stick of wood into the same bottle of gas. 
Describe the result. Is the gas combustible? Does it support com- 
bustion ? 

(c) The observations in () suggest that the gas is oxygen, but it is 
not, though this fact is not easily proved by a single experiment. Put 

Compounds of Nitrogen. 497 

a small piece of sulphur in a deflagrating spoon, light it, and lower the 
burning sulphur at once into another bottle of gas. If the experiment 
is conducted properly, the sulphur will not burn so brightly as it would 
in a bottle of oxygen. 

(cT) Stand the other bottle mouth downward in the pneumatic 
trough, or better, in a vessel of cold water. Describe the result. If 
the result is not conclusive, fill the bottle half full of water, cover with 
the hand, and shake. Does this observation help distinguish the gas 
from oxygen? 

What in all probability is the other product (seen in B) of the chemi- 
cal change in this experiment? Could it have been an impurity in the 
ammonium nitrate? What are the fumes noticed in A? 

How would you distinguish ammonium nitrate from all other nitrates ? 
How would you distinguish nitrous oxide from (#) the other oxides of 
nitrogen, (6) air, (c) oxygen, {d} hydrogen, (e) nitrogen, (/) carbon 

Experiment 57. Preparation and Properties of Sodium Nitrite. 

Materials: 10 grams sodium nitrate, 20 grams lead, iron sand bath 
pan, glass rod. 

Heat the mixture of lead and sodium nitrate on the sand bath pan, 
which stands on the ring of an iron stand. Stir the melted mass with 
a glass rod. Some of the lead will disappear and a yellowish brown 
powder will be seen in the molten mass. The action should proceed 
until most of the lead has disappeared. Allow the mass to cool, trans- 
fer to a mortar, pulverize, add hot water, and filter the clearer portion ; 
add more hot water to the residue, and filter this portion. This oper- 
ation extracts the sodium nitrite. Add to the combined filtrates several 
drops of concentrated sulphuric acid. Describe the result. How does 
the result compare with the action of concentrated sulphuric acid on 
sodium nitrate? The yellowish product is lead oxide. What general 
chemical change led to its formation? How must the nitrate have been 
changed ? 

Experiment 58. Aqua Regia. Materials: Gold leaf, concentrated 
nitric and hydrochloric acids, glass rod. 

Touch a small piece of gold leaf with the end of a moist glass rod, 
and wash the gold leaf into a test tube by pouring a few cubic centi- 
meters of concentrated hydrochloric acid down the rod. Heat gently 
until the acid just begins to boil. Does the gold dissolve? Wash an- 

498 Experiments. 

other piece of gold leaf into another test tube with concentrated nitric 
acid, and heat as before. Does the gold dissolve? Pour the contents 
of one tube into the other, and warm gently. Does the gold dissolve? 
Draw a conclusion. 


(1) What is the literal meaning and significance of the term aqua 
regia ? 

(2) What other metals does aqua regia dissolve? 

(3) What is the chemical action of aqua regia on gold? 

(4) Upon what property of nitric acid does the action of aqua regia 


Experiment 59. Distribution of Carbon. Materials: Hessian 
crucible, sand, wood, cotton, starch, sugar, glass tube (or rod), candle, 
block of wood. 

(a) Cover the bottom of a Hessian crucible with a thin layer of sand. 
Put on the sand a small piece of wood, a small, compact wad of cotton, 
and a lump of starch. Fill the crucible loosely with dry sand, and 
slip it into the ring of an iron stand. Heat with a flame which extends 
just above the bottom of the crucible until the smoking ceases (approxi- 
mately 20 minutes) . After the crucible has cooled sufficiently to handle, 
pour the contents out upon a block of wood or an iron pan. Examine 
the contents. What is the residue? What is hereby shown about the 
distribution of carbon ? 

While the crucible is heating, do the following : 

(ft) Heat about i gram of sugar in an old test tube until the vapors 
cease to appear. What is the most obvious product? 

(<:) Close the holes at the bottom of a lighted Bunsen burner, and 
hold a glass tube in the upper part of the flame long enough for a thin 
deposit to form. Examine it, name it, and state its source. 

(d) Hold a glass tube in the flame of a candle which stands on a 
block of wood, and compare the result with that in (<:). 

Draw a general conclusion regarding the distribution of carbon. 

Experiment 60. Decolorizing Action of Charcoal. Materials: 
Animal charcoal, indigo solution, filter paper and funnel. 

Fill a test tube one fourth full of powdered animal charcoal as follows : 
Fold a narrow strip of smooth paper so that it will slip easily into the 



test tube ; place the powder at one end of the troughlike holder, slowly 
push the paper into the test tube, holding both tube and paper in a 
horizontal position ; now hold the tube upright, and the powder will 
slip from the paper. Add 10 cubic centimeters of indigo solution, shake 
thoroughly for a minute, and then warm gently. Filter through a wet 
filter paper into a clean test tube. Compare the color of the filtrate 
with that of the indigo solution. Explain the change in color. 

Other organic substances besides indigo are similarly changed. 
Draw a general conclusion regarding the decolorizing power of char- 

Experiment 61. Deodorizing Action of Charcoal. Materials: 
Wood charcoal, hydrogen sulphide solution, test tube, and cork. 

Smell of a weak solution of hydrogen sulphide gas. Fill a test tube 
half full of powdered wood charcoal as in Experiment 60, add a little 
hydrogen sulphide solution, and cork securely. If the tube leaks, make 
the opening gas-tight with vaseline. Shake thoroughly. After fifteen 
or twenty minutes, remove the stopper and smell of the contents. Is the 
odor much less offensive? Repeat, unless a definite result is obtained. 
Explain the change. 

Experiment 62. Preparation of Carbon Dioxide. Materials: 
Lumps of marble, sand, concentrated hydrochloric acid, stick of wood, 
candle fastened to a wire, limewater, four bottles. Use the same 
apparatus as in the preparation of hydrogen (see Exp. 8). 

Cover the bottom of the test tube with sand, add a little water, and 
carefully slip into it half a dozen small lumps of marble. Arrange the 
apparatus to collect the gas over water, as previously directed. Add 
through the safety tube just enough concentrated hydrochloric acid to 
cover the marble. Collect four bottles, cover with glass plates or wet 
filter paper, and stand aside till needed. 

Allow the action in the flask to continue, and preserve the contents 
for Experiment 64. 

Proceed at once to the next experiment. 

Experiment 63. Properties of Carbon Dioxide. 

Study the properties of carbon dioxide gas as follows : 
(a) Plunge a burning stick several times into one bottle. Describe 
the result. 

500 Experiments. 

(b} Lower a lighted candle into a bottle of air, and invert a bottle 
of carbon dioxide over it, holding the bottles mouth to mouth. Describe 
the result. What does this result show about the density of carbon 

(V) Pour a little limewater into a bottle of carbon dioxide, cover 
with the hand, and shake vigorously. Describe and explain the 

(d) Fill a bottle of carbon dioxide one third full of water, cover 
tightly with the hand, and shake vigorously. Invert, still covered, in 
the pneumatic trough. Does the result reveal any facts about the 
solubility of carbon dioxide ? 


(1) Describe the preparation of carbon dioxide. 

(2) What do (a) and () show about the relation of carbon dioxide 
to combustion? 

(3) What is the test for carbon dioxide? 

(4) What chemical changes occur in the test for carbon dioxide? 

Experiment 64. The Solid Product of the Interaction of 
Calcium Carbonate and Hydrochloric Acid. 

Filter the contents of the test tube into an evaporating dish, adding 
a little warm water beforehand, if the contents are solid. Evaporate to 
dryness in the hood over a free flame as long as much liquid remains. 
As the residue approaches pasty consistency, add a little water, stand 
the dish on a gauze-covered support, and move the lighted burner 
underneath. Heat the residue until no fumes of hydrochloric acid are 
evolved. Dissolve some of the residue in distilled water and test 
portions for (a) a chloride and (<) a calcium compound (see 
Exp. 12 ($), (</)). If a calcium compound is found, confirm the obser- 
vation thus : Dip a clean, moist platinum test wire (see Int. 5 (4) ) 
into the solid residue, and hold it in the Bunsen flame. If calcium is 
present, the flame will be colored a yellowish red. 

What is the residue? Verify the conclusion by a simple experiment. 

Experiment 65. Carbon Dioxide and Combustion. Materials : 
Limewater, glass tube, candle attached to wire, stick of wood, two 

(a) Exhale through a glass tube into a test tube half full of lime- 
water. Describe and explain the result. 

Carbon. . 501 

() Lower a lighted candle into a bottle and allow it to burn for a few 
minutes. Remove the candle, pour a little limewater into the bottle, 
and shake vigorously. Describe and explain the result. 

(c) Allow a stick of wood to burn for a short time in a bottle (not 
the one used in (<)), and then proceed as in (). Describe the result. 
Does it confirm the results obtained in (a) and () ? 

(d*) Repeat (<:), using a piece of paper in place of the wood. De- 
scribe the result. Does it confirm the results obtained in (a), (), and 



1 i ) What is the source of the carbon dioxide in (a) ? 

(2) What is one of the gases escaping from chimneys? From a 
burning lamp ? 

Experiment 66. Carbonic Acid. (Teacher's Experiment.) 

Materials: Solutions of sodium hydroxide and phenolphthalein, bottle, 
and the carbon dioxide generator used in Experiment 62. 

Construct and arrange the carbon dioxide generator as in Experiment 
62. Fill the bottle nearly full of water, add a few drops of a solution of 
phenolphthalein * and just enough sodium hydroxide solution to color 
the liquid a faint magenta. Allow a slow current of carbon dioxide 
to bubble through the liquid in the bottle, until a definite change is 
produced in the absorbing liquid. Describe and explain it. 

Experiment 67. Preparation and Properties of Carbonates. 

Materials: Marble, sand, concentrated hydrochloric acid, limewater. 
The apparatus is the same as that used in Experiment 62. 

(a) Prepare a carbon dioxide generator as in Exp. 62, and attach it by 
a clamp to an iron stand so that the end of the delivery tube reaches to 
the bottom of a bottle half full of limewater. Pass the gas slowly into 
the limewater until considerable precipitate is formed. Remove the 
bottle and let the precipitate settle. 

(b) Meanwhile pass carbon dioxide slowly for about five minutes into 
a test tube nearly full of a dilute solution of sodium hydroxide. 

(<:) Examine the precipitate from (a) as follows : Pour off most of 
the liquid without disturbing the solid (see Int. 6 (i) (#)). Dip a glass 
tube into limewater, remove it, and a drop will adhere to the end. 

iThis compound is magenta in alkaline solutions and colorless in acid solutions. 


Pour a little hydrochloric acid into the test tube, shake, and hold the 
"limewater tube " in the escaping gas. Observe the change in the drop 
of limewater. If no change occurs, add more acid to the precipitate. 
What is the liberated gas? What is the precipitate? How was the 
latter formed? 

(d) Proceed as in (c) with the solution obtained in (V). What is 
the liberated gas? From what compound did it come? How was this 
compound formed? How does it differ from the one formed in (a) ? 


(1) What is the test for a carbonate? 

(2) How may limewater be distinguished from a solution of sodium 
or potassium hydroxide ? 

Experiment 68. Detection of Carbonates. Materials : Hydro- 
chloric acid, limewater, glass tube ; baking soda, washing soda, baking 
powder, native chalk, tooth powder, white lead, whiting, old mortar (or 

Put a little of each of the above solids in separate test tubes, add a 
little water and dilute hydrochloric acid, and shake ; hold the " lime- 
water tube " in the escaping gas, as in Experiment 67. If the action is 
not marked, warm the test tube. Describe the result in each case. 

Experiment 69. Acid Calcium Carbonate. Materials : Lime- 
water and the carbon dioxide generator used in Experiment 62. 

Pass carbon dioxide into a test tube half full of limewater until the 
precipitate disappears. Filter, if the liquid is not perfectly clear, and 
then heat. Describe the change. Explain the three changes which 
take place in the test tube. 

Experiment 70. Preparation and Properties of Carbon Mo- 
noxide. (Teacher's Experiment.) Materials : Oxalic acid, concen- 
trated sulphuric acid, limewater, pneumatic trough filled as usual, three 
bottles, three glass plates. The apparatus is shown in Figure 107 
(p. 504). 

Precaution. Carbon monoxide and oxalic acid are poisonous. Hot 
sulphuric acid is dangerous. Perform this experiment with unusual 

Put 10 grams of oxalic acid in the large test tube, A, and add 25 
cubic centimeters of concentrated sulphuric acid. Put enough lime- 
water in B to cover the end of the tube, E. The end of H should rest 

Carbon. 503 

on the bottom of the pneumatic trough just beneath the hole in the 
shelf. Heat the tube, A, gently, and carbon monoxide will be evolved. 
A small flame must be used, because the gas is rapidly evolved as the 
heat increases. It is advisable to remove or lower the flame as bubbles 
appear in the tube, B, regulate the heat by the effervescence. Collect 
all the gas, but do not use the first bottle, covering the bottles with 
glass plates as they are filled, and setting them aside temporarily. 
When the last bottle has been collected and covered, loosen the stop- 
per in B, remove the end of H from the water in the trough, and if gas 
is still being evolved, stand the whole apparatus in the hood. 
Test the gas thus : 

(a) Notice that it is colorless. 

(b) Hold a lighted match at the mouth of a bottle for an instant. 
Note the flame, especially its color and how it burns. After the flame 
has disappeared, drop a lighted match into the bottle. Describe the 
result. Draw a conclusion and verify it by (<:). 

(c} Burn another bottle of gas, and after the flame has disappeared 
pour limewater into the bottle and shake. Describe the result. 


(1) What gas besides carbon monoxide was produced, as shown 

(2) Summarize the observed properties of carbon monoxide. 

(3) What is the chemical relation of the two oxides of carbon? 

(4) How can the two oxides be changed into each other? What 
two general processes do the changes illustrate ? 

Experiment 71. Preparation and Properties of Ethylene. 
(Teacher's Experiment.) Materials: Alcohol, concentrated sulphuric 
acid, sand, pneumatic trough filled as usual, two bottles, limewater. 
The apparatus is that used in Experiment 70. f 

Precaution. A mixture of ethylene and air explodes, if ignited. 
Hot sulphuric acid is dangerous. Guard against flames, leaks, and 

Put 5 cubic centimeters of water in a test tube and slowly pour upon 
it 15 cubic centimeters of concentrated sulphuric acid. Cool the acid 
by holding the test tube in a stream of cold water. Put 5 to 7 cubic 
centimeters of alcohol in the test tube, A, add a little clean sand, and 



then slowly pour in the cold acid. The test tube, B, remains empty. A 
dish should stand under A to catch the contents, in case of accident. 

Adjust the apparatus as shown in Figure 
107, taking care not to crush the test tubes. 
Heat the test tube, A, gently between the 
bottom and the surface of the contents , to 
detect any leaks in the apparatus. Readjust, 
if necessary. Heat gently to drive out the 
air, and when it is judged that the gas which 
is being evolved is ethylene, collect two 
bottles. As the heat increases, the mixture 
is apt to froth or 
"bump"; sometimes 
the gas is evolved sud- 
denly. Hence the heat 
must be so regulated 
that the evolution of 
gas is slow. Especial 

FIG. 107. Apparatus for preparing ethylene. 

care must be taken not 
to heat the test tube 
above the surface of 
the contents, otherwise 

a sudden movement of the hot liquid might crack the test tube. As soon 
as the gas has been collected, remove the tube, //, from the water, and 
if the ethylene is still being evolved, stand the apparatus in the hood. 
When the tube, A, is cool enough to handle, pour the contents down 
the sink or into a receptacle especially provided for dangerous mixtures. 
Test the gas by holding a lighted match at the mouth of a bottle. 
Observe and record the color and temperature of the flame, its luminos- 
ity, rapidity of combustion, visible products, and any other character- 
istic properties. Repeat with the other bottle, and carefully observe 
properties needing confirmation. Add a little limewater to one of the 
bottles in which the gas was burned, shake, and explain the result. 
What evidence does this experiment present regarding the composition 
of ethylene? 

Experiment 72. Preparation and Properties of Acetylene. 

Fill a test tube nearly full of water, stand the test tube in a rack, and 
drop two or three very small pieces of calcium carbide into the test 



tube. Acetylene is evolved. After the action has proceeded long 
enough to expel the air, light the gas by holding a lighted match at the 
mouth of the tube. Observe and record the nature of the flame, espe- 
cially its color, intensity, visible products (if any), temperature, etc. 
Hold a cold glass plate or bottle over the flame. What does the result 
suggest about the composition of acetylene? What other evidence of 
its composition is revealed by the properties of the flame ? 

Experiment 73. Preparation and Properties of Illuminating 
(Coal) Gas. (Teacher's Experiment.) Materials: Soft coal, asbestos, 
pneumatic trough filled as usual, three bottles, litmus paper, filter paper, 
lead acetate (or nitrate) solution. The apparatus is shown in Figure 108. 
A A' is an ignition tube from 10 to 15 centimeters (4 to 6 inches) long. A 
spiral of copper wire is placed near A', and the tube is supported by a 

FlG. 108. Apparatus for preparing illuminating gas from soft coal. 

clamp between the wire and the end of the tube. An empty test tube 
or bottle is connected with the combustion tube by a bent tube passing 
to the bottom of B ; this vessel retains tarry matter, which comes from 
the ignition tube. The U-tube contains moistened pink litmus paper 
in the limb C, and a narrow strip of filter paper moistened with a lead 
compound (nitrate or acetate) in the limb C', the latter serving to detect 
hydrogen sulphide. The bottle, D, which maybe ar>y convenient size, 
is connected as shown in the figure, and is to be one third full of lime- 
water. The tube, ZT, is to be connected with a delivery tube passing into 
a pneumatic trough arranged to collect a gas over water. 

Fill AA' two thirds full of coarsely powdered soft coal, which should 
be held in place with a loose plug of shredded asbestos. See that all 
connections are gas-tight by heating the ignition tube gently ; if the 
ipparatus is tight, the expanded air will bubble through the bottle D. 
Readjust, if necessary. 

506 Experiments. 

Heat the whole ignition tube gently at first, and gradually increase 
the heat, but avoid heating either end very hot, otherwise the closed 
end may soften and burst or the rubber stopper may melt. As the heat 
increases, watch for marked changes in B, CC', and D. As soon as the 
slow bubbling shows that all air has been driven out of the apparatus, 
collect, as previously directed, two bottles of the gas evolved. Cover 
the bottles with wet filter paper as soon as they are removed from the 
trough. When the last bottle has been removed, disconnect the ap- 
paratus at any convenient point between A' and C. Let the ignition 
tube cool. 

Test the gas by holding a lighted match near the mouth of a bottle. 
Observe and record the color and heat of the flame. Is smoke formed ? 
Repeat with the remaining bottle, and observe more closely any facts 
suggested, but not clearly shown, by the first observations. 

Examine the contents of the ignition tube. Does it resemble coke 
or some form of carbon ? Examine the bottle, B, for tarry matter. 
Does the paper in C show the formation of ammonia ? If the paper in 
C is black or brown, it is caused by lead sulphide, which is formed by 
the interaction of hydrogen sulphide and a lead compound. Did the 
gas contain hydrogen sulphide ? Did the bottle, Z), show the formation 
of carbon dioxide ? 


(1) Describe briefly the whole experiment. 

(2) Sketch the apparatus. 

(3) Summarize the properties of coal gas. 

Experiment 74. Combustion of Illuminating Gas. Materials : 
Pointed glass tube (see Int. 3 (V)), bottle, limewater. 

Attach a pointed glass tube to the rubber tube connected with the gas 
jet, and lower a small flame into a cold, dry bottle. Observe at once 
the most definite result inside the bottle. Remove and extinguish the 
flame, add a little limewater to the bottle, and shake. What are the 
two products of the combustion of coal gas ? What do the observa- 
tions show about the composition of the main constituents of coal gas ? 

Experiment 75. Construction of a Bunsen Burner. 

Take apart a Bunsen burner and study the construction. Write 
a short description of the burner. Sketch the essential parts. 



Experiment 76. Bunsen Burner Flame. Materials: Glass tube, 
powdered wood charcoal, pin, copper wire, wire gauze. 

I. (a) Close the holes at the bottom of a Bunsen burner and hold a 
glass tube in the upper part of the flame. Note the black deposit. 
What is it? Where did it come from? Open the holes and hold the 
'blackened tube in the colorless flame. What becomes of the deposit? 
How is the flame changed, if at all? What does the experiment 
suggest about the luminosity of flame ? 

(b) Dip a glass tube a short distance into powdered wood charcoal, 
place the end containing the charcoal in one of the holes at the bottom 
of the burner, and blow gently two or three times into the other end. 
Describe and explain the result. Does it verify the answer to the last 
question in (a) ? 

(^) Open and close the holes of a lighted burner several times. 
Describe the result. Pinch the rubber tube to extinguish the flame, 
then light the gas at the holes. What change is produced in the flame? 
What causes the change? 


(1) What is the object of the holes? 

(2) Why does the gas burn at the top and not inside of the burner? 

(3) Why does the flame sometimes "strike back" and burn inside? 

(4) Why is the Bunsen flame nonluminous? 

II. (a) Hold a match across the top of the tube of 
a lighted Bunsen burner. When it begins to burn, 
Note where it is charred, 
and explain the result, 
a piece of wire 
down upon the 

remove and extinguish it. 




flame. Describe the 

appearance of the gauze. 

The same fact may be 
shown by sticking a pin through a (sul- 
phur) match, suspending it across the 
burner, and then lighting the gas. The 
position of the match is shown in Figure 

log. Turn on a full 

FIG. no. Bent tube for ex- . , . 

amining the structure of a Bun- current of g as before 
sen flame. lighting it. What does the whole 

FIG. 109. Sul- 
phur match sus- 
pended across the 
top of a Bunsen 

508 Experiments. 

experiment show about the structure of the lower part of the Bunsen 
flame? Verify your answer by (b). 

(&) Bend a glass tube about 15 centimeters (6 inches) long into 
the shape shown in Figure no. Hold the shorter arm in the flame 
about 2 centimeters (i inch) from the top of the burner tube. Hold 
a lighted match for an instant at the upper end of the tube. What' 
does the result show about the structure of^he Bunsen flame? Does it 
verify (a) ? 

(c) Find the hottest part of the flame, when a full current of gas is 
burning, by holding a copper wire in the flame. Measure its distance, 
approximately, from the top of the burner tube. 

(d?) Examine a typical Bunsen flame one which shows clearly the 
outlines of the inner part. What is the general shape of each main 
part? Draw a vertical and a cross section of the flame. 

Experiment 77. Candle Flame. Materials : Candle, two blocks 
of wood, bottle, piece of stiff white paper, limewater, matches, lamp 
chimney, copper wire ( 15 centimeters or 6 inches long). 

Attach a candle to a block of wood by means of a little melted candle 
wax, and proceed as follows : 

(a) Hold a cold, dry bottle over the lighted candle. Describe the 
result produced inside the bottle. What is the product? What is its 
source? Remove the bottle, pour a little limewater into it, and shake. 
Describe and explain the result. What are the two main products of a 
burning candle? 

(b} Blow out the candle flame, and immediately hold a lighted match 
in the escaping smoke. Does the candle relight? Why? What is the 
general nature of this smoke? How is it related to the candle wax? 
How does (b} contribute to the explanation of (a) ? 

(c) Press a piece of stiff white paper for an instant down upon the 
candle flame almost to the wick. Repeat several times with different 
parts of the paper. What does the paper show about the structure of 
the flame? 

(d) Stand a lamp chimney over the lighted candle. How is the 
flame effected? Hold the chimney a short distance (i centimeter or 
.5 inch) above the block. Does the candle continue to burn? Why? 
Keep the chimney in the same position and cover the top with a block 
of wood. What is the result ? Why ? 

Carbon. 509 

(V) Roll one end of the copper wire around a lead pencil to form a 
spiral about (2 centimeters or I inch) long. Press the spiral down 
upon the candle flame. What is the result? Why ? 


(1) Draw a candle flame, showing the parts. 

(2) What is the essential difference between a candle flame and a 
Bunsen flame ? 

(3) Is there any essential difference between a candle flame and a 
gas or a lamp flame ? 

(4) Why do candles and lamps often smoke? 

Experiment 78. Kindling Temperature. 

() Press a wire gauze down upon a Bunsen flame. Where is the 
flame ? Hold a lighted match just above the gauze. Now where is 
the flame ? 

(<$) Extinguish the flame. Turn on the gas, hold the gauze in the 
escaping gas, about 5 centimeters (2 inches) above the top of the burner, 
and thrust a lighted match into the gas above the gauze. Where is the 
flame ? Lower the gauze slowly and describe the final result. 

(c} Hold the gauze in the flame in one position for a minute or two. 
Where is the flame at the end of this time ? Why ? 


(1) Define kindling temperature. 

(2) What application is made of the principle illustrated by this 
experiment ? 

(3) State exactly how this experiment illustrates kindling tempera- 

Experiment 79. Reduction and Oxidation with the Blow- 
pipe. Materials : Blowpipe, blowpipe tube, charcoal, lead oxide 
(litharge), sodium carbonate, sodium sulphate, wood charcoal, silver 
coin, zinc, lead, tin. 

Slip the blowpipe tube into the burner, light the gas and lower the 
flame until it is about 4 centimeters (1.5 inches) high. Rest the tip of the 
blowpipe on the top of the tube, placing the tip just within the flame. 
Put the other end of the blowpipe between the lips, puff out the cheeks, 
inhale through the nose, and exhale into the tube, using the cheeks some- 
what as a bellows. Do not blow in puffs, but produce a continuous flow 
of air by steady and easy inhaling and exhaling. The operation is nat- 


ural and simple, and, if properly performed, will not make one out of 
breath. The flame should be an inner blue cone surrounded by an outer 
and almost invisible cone, though its shape varies with the method of 
production (see Fig. 44). Practice until the flame is produced volun- 
tarily and without exhaustion. Watch the flame and learn to distin- 
guish the two parts, so that they may be intelligently utilized. 

I. Reduction, (a) Make a shallow hole at one end of the flat side of 
a piece of charcoal. Fill the hole with a mixture of equal parts of pow- 
dered sodium carbonate and lead oxide, and heat the mixture in the 
reducing flame. The sodium carbonate melts and assists the fusion 
of the oxide, but the former is not changed chemically. In a short time 
bright, silvery globules will appear on the charcoal. Let the mass cool, 
and pick out the largest globules. Put one or two in a mortar, and strike 
with a pestle. Are they soft and malleable, or brittle and hard ? State 
the result when a globule is drawn across or rubbed upon a white paper. 
How do the properties compare with those of metallic lead ? What has 
become of the oxygen ? Of what chemical use is the charcoal ? 

(b) Grind together in a mortar a little sodium sulphate and wood 
charcoal, adding at intervals just enough water to hold the mass to- 
gether. Heat this paste fora few minutes in the reducing flame as in 
(a) . Scrape the fused mass into a test tube, boil in a little water, and 
put a drop of the solution on a bright silver coin. If a dark brown stain 
is produced, it is evidence of the formation of silver sulphide. Repeat, 
if no such stain is produced. State all the chemical changes which led 
to the production of the silver sulphide, explaining at the same time 
how the experiment illustrates reduction. 

II. Oxidation, (a) Heat a small piece of zinc on charcoal in the 
oxidizing flame. What is the product ? Observe its color, and the color 
of the coating on the charcoal when hot and cold. Record as described 

(b) Heat a piece of lead as in (a}. Observe the presence or absence 
of fumes, as well as the color of the coating when hot and cold. See (d). 

(c) Heat a small piece of tin in the oxidizing flame. Observe as in (b} . 
(d} Tabulate the above observations, stating (i) the color of the 

hot and cold coating on the charcoal, (2) presence or absence of fumes, 
(3) name of product. 


(1) Sketch a blowpipe. 

(2) Sketch a flame showing the oxidizing and reducing parts. 

Fluorine, Bromine, and Iodine. 511 


Experiment 80. Preparation and Properties of Hydrofluoric 

Acid. Materials : Lead dish, glass plate, paraffin, file, calcium fluoride, 
concentrated sulphuric acid. 

Precaution. Hydrofluoric acid gas is a corrosive poison. An aque- 
ous solution of the gas commercial hydrofluoric acid burns the flesh 

Warm a glass plate about 10 centimeters (4 inches) square by dipping 
it into hot water or by standing it near a warm object, such as a radiator. 
If it is held over a flame, it is liable to crack. Coat one surface with 
paraffin. The surface should be uniformly covered with a thin layer. 
Scratch letters, figures, or a diagram through the wax with a file. Be 
sure the instrument removes the wax through to the glass, and that the 
lines are not too fine. 

Put 5 grams of calcium fluoride in a lead dish and add just enough 
concentrated sulphuric acid to form a thin paste. Stir the mixture with 
a file. Place the glass plate, wax side down, upon the lead dish and 
stand the whole apparatus in the hood for several hours, or until some 
convenient time. Remove the plate. Scrape the contents of the dish, 
immediately, into a waste jar in the hood, and wash the dish free from 
acid. Most of the wax can be scraped from the glass plate with a knife. 
The last portions can be removed by rubbing with a cloth moistened 
with alcohol or turpentine. Do not attempt to melt off 
the wax over the flame. If the experiment has been 
properly performed, the plate will be etched where the 
glass was exposed to the hydrofluoric acid gas. 

Experiment 81. Preparation and Properties of 
Bromine. Materials : Potassium bromide, manganese 
dioxide, dilute sulphuric acid, bottle of water, test-tube 
holder. The apparatus is shown in Figure 1 1 1 . The 
large test tube is provided with a one-hole rubber 
stopper to which is fitted the bent glass tube. The 
latter is about 30 centimeters (12 inches) long, and is 
bent according to the directions given in the Introduc- paratus fo r pre . 
tion, 3 (b). paring bromine. 

512 Experiments. 

Precaution. Bromine is a corrosive liquid which forms, at the 
ordinary temperature, a suffocating vapor. Perform in the hood all 
experiments which use or evolve bromine. - 

Put a dozen crystals of potassium bromide in the test tube, add an 
equal quantity of manganese dioxide and 10 cubic centimeters of dilute 
sulphuric acid. Insert the stopper and its tube securely, and boil 
gently. Do not hold the test tube in the hand, but use the test tube 
holder. Brown fumes soon appear in the test tube and pass out of the 
delivery tube. Regulate the heating so that this vapor will condense 
and collect in the lower bend of the delivery tube. Both vapor and 
liquid are bromine. When no further boiling produces bromine vapor 
in the test tube, pour the bromine from the delivery tube into a bottle 
of water. Observe and record the physical properties of this bromine, 
especially the color, solubility in water, specific gravity, volatility, and 
physical state. Try the action of the contents of the bottle on litmus 
paper ; if the action is not marked, push the paper down near the bro- 
mine. Determine the odor by smelling cautiously of the water in the 
bottle. As soon as these observations have been made, pour the con- 
tents of the bottle into the sink and flush with water, or pour into a jar 
in the hood. Wash the test tube free from all traces of bromine, taking 
care to get none on the hands. 


(1) In what ways does bromine physically resemble chlorine? In 
what ways does it differ from chlorine? 

(2) How is it essentially different from all other elements previously 
studied ? 

Experiment 82. Properties of Potassium Bromide. Materials : 
Potassium bromide, silver nitrate solution, ammonium hydroxide. 

Examine a crystal of potassium bromide, and state its most obvious 
properties. Dissolve it in a test tube half full of water, and add a few 
drops of silver nitrate solution. Describe the result. Is the solid prod- 
uct soluble in ammonium hydroxide? How can bromides be distin- 
guished from chlorides ? Do the properties of bromides, typified by 
potassium bromide, suggest any marked relation to chlorides ? 

Experiment 83. Preparation and Properties of Iodine. Ma- 
terials: Potassium iodide, manganese dioxide, mortar and pestle, con- 
centrated sulphuric acid, funnel, cotton. 

Fluorine, Bromine, and Iodine. 513 

Grind together in a mortar a dozen large crystals of potassium iodide 
and about twice the bulk of manganese dioxide. Put the mixture in a 
test tube provided with a holder, moisten with water, and add a few 
cubic centimeters of concentrated sulphuric acid. Plug with cotton the 
inside opening of a funnel, and hold the latter firmly over the mouth of 
the test tube. Heat the test tube gently with a low flame (5 centime- 
ters or 2 inches). The vapor of iodine will fill the test tube, and crys- 
tals will collect in the upper part of the test tube and in the funnel. 
If the crystals collect in the test tube, a gentle heat will force them 
into the funnel. Continue to heat until enough iodine collects in the 
funnel for several experiments. Scrape the crystals into a dish. 

Study the properties as follows : 

() Observe and record the physical properties of iodine, especially 
the color of the solid and of the vapor, volatility, and odor (cautiously). 

() Heat a crystal in a dry test tube, and when the tube is half full 
of vapor, invert it. What does the result show about the density of 
iodine vapor? 

(c) Touch a crystal with the finger. What color is the stain ? Will 
water remove it? Will alcohol? Will a solution of potassium iodide? 
What do these results show about the solubility of iodine ? 

(NOTE. If crystals are left, use them in the next experiment. Pre- 
serve in a stoppered bottle.) 

Experiment 84. Test for Iodine with Carbon Bisulphide. 
Materials : Iodine, potassium iodide, carbon disulphide, chlorine water. 

Precaution. Carbon disulphide is inflammable. It should not be 
used near flames. 

(a) Free iodine. Add a few drops of carbon disulphide to a very 
dilute solution of iodine, made by dissolving a crystal of iodine in a 
solution of potassium iodide, and observe the color of the carbon disul- 
phide, which, being much heavier than water, will sink to the bottom of 
the test tube. How does it resemble the color of iodine vapor ? 

(b) Combined iodine. Add a few drops of carbon disulphide to a 
very dilute solution of potassium iodide. Is there positive evidence of 
iodine ? Now add several drops of chlorine water, and shake. How 
does this result compare with the final result in (a) ? The result is due 
to the fact that chlorine liberates iodine from its compounds, and the 
iodine, being free, exhibits the characteristic color. 


Experiment 85. Test for Iodine with Starch. Materials: 
Starch, mortar and pestle, iodine solution, potassium iodide, chlorine 

Grind a lump of starch in a mortar with a little water to creamy con- 
sistency. Pour this into about 100 cubic centimeters of boiling water, 
and stir the hot liquid. Allow it to cool, or cool it by holding the 
vessel in a stream of cold water, and then pour off the clear liquid. 
Use this cold starch solution to test for iodine. 

(a) Free iodine. Add a few cubic centimeters of the starch solution 
to a test tube nearly full of water, and then add a few drops of iodine 
solution. The deep blue color is due to the presence of a compound 
which is always formed under these circumstances, but the composition 
of which is unknown. If the color is black, pour out half of the liquid 
and add more water, or pour some of the liquid into a dish of water. 

(b) Combined iodine. Add a few cubic centimeters of the starch 
solution to a very dilute solution of potassium iodide. Is the blue com- 
pound formed ? Add a few drops of chlorine water, and shake. Com- 
pare with the final result in Experiment 84 (b) . 

Experiment 86. Detection of Starch by Iodine. Materials: 
Dilute solution of iodine (in potassium iodide), mortar and pestle, 
potato, rice, bread. 

Test the potato, rice, and bread for starch by grinding a little of each 
with water in a mortar, and then adding a few drops of the extract to a 
very dilute solution of iodine. State the result in each case. 

Experiment 87. Properties of Potassium Iodide. Materials : 
Potassium iodide, silver nitrate solution, ammonium hydroxide. 

Proceed with the potassium iodide as in Experiment 82. 

How can iodides be distinguished from chlorides ? Do iodides, 
typified by potassium iodide, suggest any marked relation to bromides 
and chlorides ? 


Experiment 88. Properties of Sulphur. 

(a) Examine a lump of sulphur, and state briefly its most obvious 
physical properties. 

(6) Optional. Weigh a lump of roll sulphur to a decigram. Slip it 
carefully into a graduated cylinder previously filled with water to a 

Sulphur and its Compounds. 515 

known point about half full and note the increase in the volume of 
water. This increase in volume is equal to the volume of the sulphur. 
Calculate the specific gravity of sulphur from the observed data. 

(NOTE. Specific gravity equals weight in air divided by weight of 
equal volume of water.) 

Experiment 89. Amorphous Sulphur. Materials: Sulphur, old 
test tube, evaporating dish. 

Put a few pieces of roll sulphur in an old test tube. Heat carefully 
until the sulphur boils, and then quickly pour the contents of the test 
tube into a dish of cold water. This is amorphous sulphur. Note its 
properties. Preserve, and examine it after twenty -four hours. Describe 
the change, if any. 

Define amorphous, and illustrate it by this experiment. 

Experiment 90. Crystallized Sulphur. J Materials : Sulphur 
(roll and flowers), Hessian crucible, carbon disulphide, evaporating 

(a) Monodinic. Fill a small Hessian crucible nearly full of roll 
sulphur. Support the crucible in the ring of an iron stand, and heat 
until all the sulphur is melted. Let it cool, and as soon as crystals 
shoot out from the walls just below the surface, pour the remaining 
melted sulphur into a dish of cold water. When the crucible can be 
handled without discomfort, crack it open lengthwise. Observe and 
record the properties of the crystals, especially the shape, size, color, 
luster, brittleness, and any other characteristic property. Allow the 
best crystals to remain undisturbed for a day or two ; then examine 
again, and record any marked changes. 

(b} Orthorhombic. Put 3 grams of flowers of sulphur in a test tube 
and add about 5 cubic centimeters of carbon disulphide remember 
the precaution to be observed in using this liquid (see Exp. 84). 
Shake until all the sulphur is dissolved, then pour the clear solution 
into an evaporating dish to crystallize. It is advisable, though not 
absolutely necessary, to stand the dish in the hood or out of doors, 
where there is no flame and where the offensive vapor will be quickly 
removed. Watch the crystallization toward the end, and, if perfect 
crystals form, remove them with the forceps (see Fig. 49). Allow the 

i See Appendix, 3 (3), (5). 


liquid to evaporate almost entirely, then remove and dry the crystals. 
Examine them as in (a) and record their properties. 

(1) Tabulate the essential results in (a) and (). 

(2) Make an outline sketch of an orthorhombic crystal of sulphur. 

Experiment 91. Combining Power of Sulphur. Materials : 
Sulphur, deflagrating spoon, bottle, iron powder, hydrochloric acid. 

(a) Set fire to a little sulphur in a deflagrating spoon, and lower 
the spoon into a bottle. Cautiously waft the fumes toward the nose, 
and observe and describe the odor. The product is a mixture of two 
oxides of sulphur. What does their formation show about the combin- 
ing power of sulphur ? 

() Repeat Experiment 34. 

Results similar to that in (<) are obtained with copper and other 
metals. Draw a general conclusion regarding the power of sulphur to 
combine with metals. 

Experiment 92. Sulphur and Matches. 

(a) Examine a sulphur match. Do you detect any sulphur ? Where? 

(b) Light a sulphur match, and observe the entire action, as far as 
the sulphur is concerned. Describe it. 

(c) What is the function of the sulphur in a burning match ? 

Experiment 93. Preparation of Hydrogen Sulphide. Mate- 
rials: Ferrous sulphide, dilute hydrochloric acid, three bottles, three 
glass plates, stoppered bottle, litmus paper. Use the same apparatus 
as in Experiment 38. 

Precaution. Hydrogen sulphide is a poisonous gas and has an 
offensive odor. It should not be inhaled. Perform in the hood all ex- 
periments evolving hydrogen sulphide. 

(a) Construct and arrange an apparatus like that shown in Figure 104. 
Fill the test tube, A, one third full of coarsely powdered ferrous sulphide, 
insert the stopper tightly, pour enough hydrochloric acid through the 
safety tube to cover the contents of the test tube. Hydrogen sulphide 
gas is rapidly evolved. If the evolution of gas slackens or stops, warm 
gently or add more hydrochloric acid. Collect three bottles, removing 
each as soon as full and covering with a glass plate. Set aside until 

Sulphur and its Compounds. 517 

() As soon as the last bottle of gas has been removed and covered, 
put in its place a bottle one fourth full of water. Adjust its height (by 
wooden blocks or by lowering the generator) so that the end of the 
delivery tube reaches to the bottom of the bottle. Continue to pass 
the gas into the water, by heating the test tube if necessary. The gas 
will be absorbed by the water, forming hydrogen sulphide water. 
Preserve it in a stoppered bottle for Experiment 95. 

Proceed at once with next experiment. 

Experiment 94. Properties of Hydrogen Sulphide Gas. 

Study as follows the hydrogen sulphide gas prepared in Experi- 
ment 93 : 

(#) Waft a little of the gas cautiously toward the nose, and describe 
the odor. This is characteristic of hydrogen sulphide, and is a decisive 
test. Has the gas color? 

(b) Test the gas from the same bottle with both kinds of moist 
litmus paper. Is it acid, alkaline, or neutral ? 

(c) Bring a lighted match to the mouth of the same bottle. Observe 
the properties of the flame as in previous experiments. Observe cau- 
tiously the odor of the product of the burned gas ; to what compound 
is the odor due? What, then, is one component of hydrogen sulphide ? 

(d) Burn another bottle of hydrogen sulphide and hold a cold bottle 
over the burning gas. What additional experimental evidence does this 
result give regarding the composition of hydrogen sulphide ? 

(V) Repeat any of the above with the remaining bottle of gas. 


(1) Summarize the properties of hydrogen sulphide gas. 

(2) State the experimental evidence of its composition. 

Experiment 95. Preparations and Properties of some Sul- 
phides. Materials: Hydrogen sulphide water prepared in Experiment 
93, clean copper wire, clean sheet lead, bright silver coin, lead oxide 
(litharge) ; solutions of lead nitrate, arsenic trioxide (in hydrochloric 
acid), tartar emetic, zinc sulphate. 

(a) Shake the bottle of hydrogen sulphide water prepared in Experi- 
ment 93 (or a similar solution), and hold successively at the mouth or 
in the neck of the bottle (i) a clean copper wire, (2) a bright strip 
of lead, and (3) an untarnished silver coin. Describe the result in 
each case. These compounds are sulphides of the respective metals. 

518 Experiments. 

() Put a little litharge the brownish yellow oxide of lead 
in a test tube, cover it with hydrogen sulphide water, and warm 
gently. The product is lead sulphide. Describe it. Explain the 

(c) Add hydrogen sulphide water to lead nitrate solution. The 
product is lead sulphide. Observe the color. 

(*/) Proceed as in (c) with the arsenic solution. Observe the color 
of the arsenic sulphide. 

(e) Proceed as in (c) with the tartar emetic solution. Tartar emetic 
is a compound of antimony. Observe the color of the antimony 

(_/") Proceed as in (c) with the zinc sulphate solution. Observe the 
color of the zinc sulphide. 

Experiment 96. Preparation of Sulphur Dioxide. Materials : 
Sodium sulphite, concentrated sulphuric acid, litmus paper, three 
bottles, two glass plates, stick of wood, pink flower. The apparatus is 
constructed, arranged, and used as in Experiment 41, with one excep- 
tion. The safety tube must be replaced by a 'dropping tube made thus . 
Cut off the top of a thistle tube about 2.5 centimeters (i inch) below 
the juncture of the stem and cup, slip a short rubber tube (5 centi- 
meters, 2 inches, long ) over one end of the stem, attach a Mohr ? s pinch- 
cock to the rubber tube, and connect the tube with the cup. 

(a) Put about 10 grams of sodium sulphite in the large test tube, 
cover with water, and insert the stopper with its tubes. Adjust the ap- 
paratus as shown in Figure 104. Fill the cup with concentrated sulphuric 
acid, open the pinchcock a little, and let the acid flow drop by drop upon 
the sodium sulphide. Sulphur dioxide gas is evolved and passes into 
the bottle, which should be removed when full, as previously described. 
Moist blue litmus paper held at the mouth of the bottle will show when 
the latter is full. Collect two bottles of gas, cover each with a glass 
plate, and set aside until needed. 

(b) As soon as the second bottle of gas has been removed and 
covered, put in its place a bottle one fourth full of water. Adjust its 
height (if necessary) by wooden blocks, so that the end of the delivery 
tube is just above the surface. Continue to add the acid drop by drop, 
at intervals, and the gas will be absorbed by the water. Shake the bottle 

Meanwhile study the gas already collected. 

Sulphur and its Compounds. 519 

Experiment 97. Properties of Sulphur Dioxide Gas. 

Proceed as follows with the gas prepared in Experiment 96 (#) : 

(a) Observe and state the most obvious physical properties, e.g. 
color, odor (cautiously), density. 

(b) Hold a blazing stick in a bottle of the gas. Will the gas burn 
or support combustion ? What previously acquired facts would have 
enabled you to predict this result ? 

(c) Pour water into the same bottle of sulphur dioxide until half full, 
cover with the hand, and shake. What is the evidence of solution ? 
Is the resulting liquid acid, alkaline, or neutral ? 

(d) Moisten a pink flower with a few drops of water, hang it in the 
remaining bottle of sulphur dioxide, holding it in place by putting the 
stem between the glass and a cork. Observe and describe any change 
in the color of the flower. What is this operation called ? 

Experiment 98. Properties of Sulphurous Acid. 

Test as follows the solution of sulphurous acid prepared in Experi- 
ment 96 () : 

(a) Taste cautiously, and describe the result. 

(b) Apply the litmus test, and state the result. 

(c} Pour a few drops of concentrated sulphuric acid into the bottle. 
What gas is liberated ? 

Experiment 99. Action of Sulphuric Acid with Organic Matter. 

Materials : Concentrated sulphuric acid, sheet of white paper, sugar, 
starch, stick of wood. 

(a) Write some letters or figures with dilute sulphuric acid on a sheet 
of white paper, and move the paper back and forth over a low flame, taking 
care not to set fire to the paper. As the water evaporates the dilute 
acid becomes concentrated. Observe and describe-the result. Paper is 
largely a compound of carbon, hydrogen, and oxygen, and the hydrogen 
and oxygen are present in the proportion to form water. Explain the 
general chemical change in this experiment. 

(b} Fill a test tube one fourth full of sugar, add an equal bulk of 
water, stand the test tube in the rack, and add cautiously several drops 
of concentrated sulphuric acid. If there is no decided result, add 
more acid. What is the black product ? Compare the final result with 
that obtained in Experiment 59 (). Is the chemical action the same in 

520 Experiments. 

each experiment ? Are the statements made in (a) about paper also 
true of sugar ? 

(c) Repeat (), using powdered starch instead of sugar. Describe 
the result. How does the result resemble that in (b} and in Experi- 
ment 59 (a) ? Predict the components of starch. In what simple way 
may the prediction be verified ? 

(//) Stand a stick of wood in a test tube one fourth full of concen- 
trated sulphuric acid. Allow it to remain in the acid for fifteen minutes, 
then remove the stick and wash off the acid. Describe the change in 
the stick. Does it resemble that in (#), (), and (c), and in Experiment 

Experiment 100. Test for Sulphuric Acid and Sulphates. 

Materials: Sulphuric acid, sodium sulphate, barium chloride solution, 
calcium sulphate, charcoal, powdered charcoal, blowpipe, silver coin. 

(a) Repeat Experiment 12 (c) with sulphuric acid and with sodium 
sulphate solution. 

() Repeat Experiment 79 I () with calcium sulphate instead of 
sodium sulphate. 


(1) State briefly the test for sulphuric acid and soluble sulphates. 
For insoluble sulphates. 

(2) How can a sulphate be distinguished from a sulphite ? 


Experiment 101. Preparation and Properties of Silicic Acid. 

Materials: Sodium silicate solution, hydrochloric acid, evaporating 
dish, gauze-covered ring. 

Add dilute hydrochloric acid to a test tube half full of sodium silicate 
solution, and shake. The jellylike precipitate is silicic acid. Rub 
some between the fingers and describe the result. Evaporate the 
precipitate to dryness in a porcelain dish which stands upon a gauze- 
covered ring in the hood. As the mass hardens, stir it with a glass rod. 
Toward the end, add more hydrochloric acid and evaporate to complete 
dryness. Then heat strongly for five minutes. The residue is silicon 
dioxide mixed with chlorides of sodium and potassium. Rub some 
between the fingers or across a glass plate. Is any grit detected ? State 
the chemical changes which occur in changing sodium silicate into 
silicon dioxide. 

Silicon and Boron. 521 

Experiment 102. Tests with Borax Beads. Materials: Pow- 
dered borax, platinum test wire (see Int. 5 (4)), solutions of cobalt 
nitrate and copper sulphate, manganese dioxide. 

Make a small loop on the end of the platinum test wire, moisten it, 
and dip it into powdered borax. Heat it in the flame, rotating it slowly ; 
at first the borax swells, but finally shrinks to a small, transparent 
bead. If the bead is too small add more borax and heat again. After 
use, the bead may be removed by dipping it, white hot, into water ; the 
sudden cooling shatters the bead, which may then be easily rubbed or 
scraped from the wire. 

(a) Cobalt Compounds. Touch a transparent borax bead with a 
glass rod which has a drop of cobalt nitrate solution on the end. Heat 
the bead in the oxidizing flame. Observe the color when cold. If it is 
black melt a little more borax into the bead ; if faintly colored, moisten 
again with the cobalt solution. The color is readily detected by look- 
ing at the bead against a white object in a strong light, or by examining 
it with a lens. When the color has been definitely determined, heat 
again in the reducing flame. Compare the color of the cold bead with 
the previous observation. 

(b) Copper compozinds. Make another transparent bead, moisten it 
with copper sulphate solution and heat it first in the oxidizing flame, 
and then in the reducing flame. Compare the colors of the cold beads, 
and draw a conclusion. 

(c) Manganese Compounds. Make another transparent bead, touch 
it with a minute quantity of manganese dioxide, and proceed as in (b). 
Compare the colors of the cold beads, and draw a conclusion. 

(d) Tabulate the results of this experiment. 


Draw a Bunsen flame, showing the reducing and oxidizing parts. 

Experiment 103. Preparation and Properties of Boric Acid 
and the Test for Boron. Materials: Borax, alcohol, evaporating 
dish, concentrated hydrochloric acid. 

To a test tube half full of boiling water, add about 10 grams of 
powdered borax. Add about 5 cubic centimeters of concentrated 
hydrochloric acid to this hot solution, and let the whole cool. Crystals 
of boric acid will separate. Filter. Describe the crystals. 

Put some of the crystals in an evaporating dish, add a little alcohol, 

522 Experiments. 

and set fire to the solution. Observe the color of the flame. It is 
caused by a complex compound of boron, and is the test for this 

Experiment 104. Some Properties of Phosphorus. 

(a) Smell of the head of a phosphorus-tipped match. Describe the 

(b) Rub the head of a phosphorus-tipped match in a dark place, and 
observe and describe the result. 

(c) The most striking property of phosphorus is the readiness with 
which it lights and burns in air. This property is too dangerous to try 
in the laboratory. Read about it in the text book. What application 
is made of this property ? Why ? 

Experiment 105. Test for Arsenic. 

Repeat Experiment 95 (</). 

Experiment 106. Test for Antimony. 

Repeat Experiment 95 (e). 

Experiment 107. Test for Bismuth. Materials: Bismuth, aqua 

Prepare a solution of bismuth chloride by heating the metal with aqua 
regia. Fill the test tube half full of water, and describe the result. The 
product is bismuth oxychloride. How is it related to bismuth chloride ? 


Experiment 108. Properties of Sodium. Materials: Sodium, 
pneumatic trough filled with water as usual, litmus paper, filter paper, 
tea lead. 

Precaution. Observe the precautions as in Experiment 24. 

(a) Examine a small piece of sodium, and record its most obvious 
physical properties, e.g. color, luster, whether hard or soft, etc. 

() Repeat Experiment 24 (except (*)). 


(1) Is sodium heavier or lighter than water? 

(2) What properties show that it is a metal? 

Sodium. 523 

(3) Is it harder or softer than most metals ? 

(4) What is the test for sodium ? 

Experiment 109. Preparation and Properties of Sodium Hy- 
droxide. Materials : Sodium carbonate, lime, zinc sulphate solution, 
iron (or tin) dish, file. 

Dissolve 25 grams of sodium carbonate in 150 cubic centimeters of 
water and heat gently in an iron dish (an ordinary iron spider is well 
adapted for this work). Meanwhile slake 10 grams of lime by adding 
just enough water to make a milky liquid " milk of lime." Add the 
milk of lime to the sodium carbonate solution and boil for several min- 
utes, stirring constantly with a file. Let the precipitate settle, remove 
a little liquid with a small tube, and if it effervesces with hydrochloric 
acid, add more milk of lime and boil; if not, pour the liquid into a con- 
venient vessel, let it stand for a few minutes or until the solid settles ; 
then pour the liquid down a glass rod ( see Int. 6 (i) ) into a bottle. 
This solution of sodium hydroxide may be evaporated to dryness, and 
the solid product tested and the remainder preserved, or the solution 
may be tested at once as follows : 

(a) Rub a little between the fingers and describe the feeling. 

($) Apply the litmus test. Is it acid or alkaline? Is it markedly so? 

(c) Add a little to a zinc sulphate solution, and shake. The precipi- 
tate is zinc hydroxide. Describe it. Now add an excess of sodium 
hydroxide, and shake. Describe the result. The excess of sodium 
hydroxide forms soluble sodium zincate. This behavior of zinc com- 
pounds is the test for an hydroxide. 

(d} How do sodium compounds affect a colorless flame? Try it, if 
in doubt. 

Exercises for Review. 

1. How does sodium carbonate act when exposed to air? Sodium 
sulphate? Sodium hydroxide? 

2. How does sodium carbonate solution affect litmus paper? How 
does its action compare with that of other salts, sodium chloride for 
example ? 

3. What marked property has sodium carbonate, and how is this 
property utilized ? 

4. How does sodium bicarbonate interact with tartaric acid ? Would 
the action be the same with other acids? 


5. " Sodium thiosulphate forms a supersaturated solution." Explain 
and illustrate this statement. 

6. How does sodium chloride interact with sulphuric acid ? Sodium 
nitrate with sulphuric acid ? 


Experiment 110. Properties of Potassium. Materials: Potas- 
sium, pneumatic trough filled as usual, litmus paper. 
Precaution. Observe the same precaution as in using sodium. 

(a) Examine a very small piece of freshly cut potassium, and record 
its most obvious physical properties. Touch it slightly. Does it sug- 
gest caustic potash and soda? 

(b) Drop a small piece of potassium on the water in a pneumatic 
trough. Stand just near enough to see the action. Describe the 
action. How does it differ from. the action of sodium? Test the water 
as in Experiment 24 (d). 

From what has already been learned about sodium and potassium, 
predict the main chemical change observed in (). 


(1) Is potassium heavier or lighter than water? 

(2) What properties suggest that it is a metal? 

(3) How does it resemble and differ from sodium? 

(4) How does potassium color a flame ? 

(5) What is the test for potassium? 

Experiment 111. Preparation and Properties of Potassium Hy- 
droxide. Materials: Potassium carbonate, lime, zinc sulphate solu- 
tion, iron (or tin) dish, file. 

Proceed as in Experiment 109, but use potassium carbonate instead of 
sodium carbonate. Test as in the case of sodium hydroxide. 

Experiment 112. Preparation and Properties of Potassium 
Carbonate. Materials: Cream of tartar, wood ashes, litmus paper, 
hydrochloric acid, iron sand bath pan, mortar and pestle. 

(a) Heat strongly 5 grams of cream of tartar acid potassium tar- 
trate in an iron pan in the hood until the residue is nearly white. 
Grind this solid with water in a mortar, and filter. Test the filtrate 
(i) with both kinds of litmus paper, (2) for potassium, and (3) for a 
carbonate. Record the results- 

Copper. 525 

() Fill a test tube half full of wood ashes, add half the volume of 
water, shake, and warm gently. Filter, and test the filtrate as in (#). 
If test (3) is not decisive, repeat the experiment on a larger scale. 
Record the results. 

(c) Expose a little potassium carbonate to the air for an hour or more. 
Describe the result. How does its behavior compare with that of sodium 
carbonate under the same conditions ? 


(1) What is the source of cream of tartar? 

(2) What do (a) and () show about the distribution of potassium? 
Of its assimilation by plants? 

(3) What is the literal meaning of the word potash ? 

Exercises for Review. 

1. What does potassium chlorate yield when heated? 

2. Does potassium chlorate dissolve readily in cold water? In hot 
water ? 

3. What is formed by heating potassium bromide with manganese 
dioxide and sulphuric acid ? 

4. Apply question 3 to potassium iodide. 

5. What happens to potassium hydroxide when exposed to air ? To 
potassium carbonate? 

6. Of what important mixture is potassium nitrate an ingredient ? 


Experiment 113. Physical Properties of Copper. 

Examine several forms of copper wire, sheet, filings, etc. and 
state the most obvious physical properties. 



(1) Is copper a good conductor of heat ? Of electricity ? On what 
evidence is your answer based ? 

(2) Is copper ductile ? Malleable ? Brittle ? Tough ? Hard or 

(3) What happens to copper when heated ? When exposed to the 
air ? 

526 Experiments. 

Experiment 114. Tests for Copper. Materials : Copper wire, 
copper sulphate solution, ammonium hydroxide, acetic acid, potassium 
ferrocyanide solution. 

(/i) Heat a copper wire in the Bunsen flame. The color is charac- 
teristic of copper and its compounds, though not a conclusive test, 
since the same color is produced by other substances. 

(^) Add a few drops of ammonium hydroxide to copper sulphate 
solution, and observe the result; now add an excess of ammonium 
hydroxide. The blue solution is a characteristic and decisive test 
for copper. 

(c) Add to a test tube one fourth full of water an equal volume of 
copper sulphate solution, and shake ; then add a few drops of acetic 
acid and of potassium ferrocyanide solution. The brown precipitate 
is copper ferrocyanide. 

Experiment 115. Interaction of Copper with Metals. Mate- 
rials: Copper wire, iron nail, zinc, solutions of copper sulphate and 
any mercury compound. 

(a) Put a clean copper wire into a solution of any mercury compound. 
After a short time, remove the wire and wipe it with a soft cloth or 
paper. Describe the change. What has become of some of the copper? 

() Put in separate test tubes half full of copper sulphate solution 
a bright iron nail and a strip of clean zinc. After a short time remove 
the metals and examine them. What is the deposit ? What has become 
of some of the zinc and iron? Does the final color of the solution 
indicate any chemical change ? How would you prove the answer to 
the last question ? 

Exercises for Review. 

1. What happens to a crystal of copper sulphate when heated ? 

2. You are given a blue solution supposed to be copper sulphate. 
State how you would prove it. 

3. What are formed by the interaction of copper and nitric acid? 

4. What color have many copper compounfls? 

5. " Brass is an alloy of copper." State how you would prove this. 


Experiment 116. Preparation of Silver. Materials : (#) Sil- 
ver nitrate solution, mercury, evaporating dish ; (b) ten-cent piece. 

Silver and Gold. 527 

concentrated nitric acid, hydrochloric acid, sulphuric acid, zinc, evapo- 
rating dish, charcoal, sodium carbonate, blowpipe. 

Prepare silver by one or both of the following methods : 

(a} Fill a porcelain dish half full of silver nitrate solution, and add a 
few drops of mercury. Allow the whole to stand undisturbed for a day 
or more, and then examine. The delicate crystals attached to the 
mercury are silver. Pick them out with the forceps, wash well with 
water, and preserve them for Experiment 117. 

() Dissolve a ten-cent piece in 10 cubic centimeters of concentrated 
nitric acid, dilute with an equal volume of water, and add hydrochloric 
acid until the precipitation is complete. Let the precipitate settle. 
Then filter, and wash until the filtrate is neutral. If convenient, let 
the precipitate dry ; if not, scrape half from the opened paper with a 
knife, put it in a porcelain dish, cover with dilute sulphuric acid, and 
add a piece of zinc ; put the other half in a cavity at the end of a piece 
of charcoal, cover with sodium carbonate, and reduce it with a blowpipe 
flame. In the first case, the silver will collect as a grayish powder; 
remove any excess of zinc, filter, wash with water and dry the residue. 
It may be preserved as a powder, or fused into a bead with a blowpipe 
flame. In the second case, minute globules of silver will appear on the 
charcoal ; scrape them together and fuse into a single bead. Preserve it. 

Experiment 117. Properties of Silver. 

Examine the silver formed in Experiment 116, and state briefly its 
most obvious properties. 

Experiment 118. Test for Silver. 

Devise a test for combined silver, based upon pYevious experiments. 
Verify it. 

Exercises for Review. 

1. What is formed by the interaction of silver nitrate and potassium 
chloride? Potassium bromide? Potassium iodide? How do the 
products differ ? 

2. What caused the blue filtrate in Experiment 116 (b) ? What 
metal besides silver does a ten-cent piece contain ? 

3. What compound is formed when silver tarnishes? 

4. What is the test for a chloride ? 

528 Experiments. 

Experiment 119. Interaction of Gold and Aqua Regia and 
the Test for Gold. Materials : Gold leaf, concentrated nitric and 
hydrochloric acids, stannous chloride solution. 

Prepare a solution of gold chloride according to Experiment 58, using 
as small a volume of the acids as possible. Dilute with water, and then 
slowly add a dilute solution of stannous chloride. A precipitate is 
produced, varying in color from faint purple to black according to the 
conditions. This precipitate is supposed to be finely divided gold, and 
is called Purple of Cassius ; its formation is the test for gold. 


Experiment 120. Tests for Calcium. 

() Subject calcium chloride to the flame test. Record the result. 
(b) Repeat Experiment 12 (d). 

Experiment 121. The ' ' Setting ' ' of Plaster of Paris. Materials: 
Plaster of Paris, block of wood. 

Mix a little plaster of Paris with enough water on a block of wood to 
form a thin paste. Let it stand undisturbed for a few minutes, and 
then examine. Describe the change. How is this property utilized ? 

Exercises for Review. 

1. Describe calcium chloride. How does it act when exposed to 
the air ? How would you show that it (a) contains calcium and (6) is 
a chloride ? 

2. Describe lime. What effect does water have upon it ? 

3. What compounds are produced by the interaction of calcium car- 
bonate and hydrochloric acid ? 

4. What is limewater ? Milk of lime ? 

5. What is formed when carbon dioxide is passed into limewater ? 
Into sodium hydroxide ? 

6. What compounds are formed by heating calcium carbonate ? 

7. What happens when an excess of carbon dioxide is passed into 
limewater ? 

8. What is acid calcium carbonate ? How can it be changed into 
normal calcium carbonate ? 

9. What compounds are formed by the interaction of calcium fluoride 
and sulphuric acid ? 

Barium. 529 

10. What is calcium hypochlorite ? For what is it used ? 

1 1 . What happens to crystallized calcium sulphate (selenite) when 
heated ? 

12. Calcium sulphate is nearly insoluble in water; how can it be 
shown to be a sulphate ? 

13. What is tricalcium phosphate ? Superphosphate of lime ? 

14. What is the scientific name of lime, limestone, marble, gypsum, 
fluor spar, limewater, slaked lime, quicklime, bleaching powder, chloride 
of lime ? 

Experiment 122. Test for Strontium. 

Dip a platinum test wire (or a glass rod) into a solution of strontium 
nitrate, and hold it in the Bunsen flame. Describe the result, after 
several trials. 

Experiment 123. Red Fire. Materials: Strontium nitrate, pow- 
dered potassium chlorate, powdered shellac, iron pan or brick. 

Mix carefully small and equal (in bulk) quantities of the three sub- 
stances on a sheet of paper. Place the mixture on a sand-bath pan or 
a brick in the hood, and light it with a Bunsen burner. Describe the 


(Compounds of Barium are Poisonous?) 
Experiment 124. Tests for Barium. 

(a) Repeat Experiment 122, using a solution of barium chloride or 
barium nitrate. Be sure the test wire (or rod) is clean. Describe the 

(b) Devise a test. (Suggestion. What is the test for a sulphate ?) 

Experiment 125. Green Fire. 

Repeat Experiment 123, using barium nitrate instead of strontium 

Exercises for Review. 

1. State the properties of barium sulphate. 

2. How does barium chloride behave toward litmus ? 

3. What industrial use have the barium oxides ? 

530 Experiments. 


Experiment 126. Properties of Magnesium. 
Examine a piece of magnesium and state briefly its most obvious 

Experiment 127. Tests for Magnesium. Materials: Solutions 
of magnesium sulphate (or chloride), ammonium chloride, ammonium 
hydroxide, disodium phosphate, and cobaltous nitrate ; magnesium 
oxide, charcoal, blowpipe. 

(a) To a solution of magnesium sulphate (or chloride) add succes- 
sively solutions of ammonium chloride, ammonium hydroxide, and di- 
sodium phosphate. A precipitate of ammonium magnesium phosphate 
is formed. It is voluminous at first, but finally crystalline. It is soluble 
in acids. Try it. 

(b) Put a little powdered magnesium oxide in a cavity at the end 
of a piece of charcoal, moisten with water, and heat intensely in a 
blowpipe flame. Cool, and moisten with a drop of cobaltous nitrate 
solution. Heat again, and when cool observe the color. If the experi- 
ment has been conducted properly, a pink or pale flesh-colored residue 
coats the charcoal. Describe the result. 

Exercises for Review. 

1. What compound is formed by burning magnesium in air or in 
oxygen? Describe it. 

2. How was magnesium utilized in the discovery of argon ? 

Experiment 128. Properties of Zinc. 

Examine a piece of zinc and record its most obvious properties. 

1. What happens to zinc when it is heated? Describe and name 
the product. 

2. Is zinc hard or soft? Malleable? Ductile? Brittle? Tough? 
Does it melt readily? 

Experiment 129. Tests for Zinc. Materials : Zinc oxide, cobalt- 
ous nitrate solution, charcoal, blowpipe. 

Mercury. 531 

(a) Recall or devise a simple test for combined zinc. (Suggestion. 
See Exp. 109 (<:).) 

(b) Recall or repeat the action of zinc when heated in the oxidizing 
flame. (See Exp. 79 II (a).) 

(<;) Fill a small cavity at one end of a piece of charcoal with zinc 
oxide, moisten with water, and heat strongly in the blowpipe flame. 
Cool, and moisten with a drop of cobaltous nitrate solution, then heat 
again. Cool and examine. A green incrustation is caused by zinc 

Experiment 130. Interaction of Zinc and Metals. Materials : 
Sheet zinc, solutions of copper sulphate, lead nitrate, mercurous nitrate. 

(a) Repeat Experiment 115 (). 

(b) Repeat (a) using lead nitrate solution. 

(c) Repeat (#) using the mercury salt solution. Examine after a 
short time, and describe. What is amalgamated zinc, and for what is it 

Exercises for Review. 

1. What are formed by the interaction of zinc and sulphuric acid? 
Of zinc and nitric acid ? 

2. What is formed by the interaction of a zinc salt and a little 
sodium hydroxide solution ? An excess of the alkali ? 


Experiment 131. Test for Cadmium. 

Add hydrogen sulphide water to a test tube half full of cadmium 
chloride solution. The precipitate is cadmium sulphide. Describe it. 
Let it settle, pour off most of the liquid, fill the test tube half full with 
dilute sulphuric acid, and warm. Describe the result. 


{Mercury and its Compounds are Poisonous.} 
Experiment 132. Properties of Mercury. 

(a) Pour a drop or two of mercury into an evaporating dish. Ex- 
amine the mercury, and state its characteristic properties. Agitate the 
dish, and describe the result. Why is mercury called u quicksilver " ? 

(b) Lift carefully a bottle of mercury. Estimate the specific gravity. 
Verify the estimate by consulting a book. 

532 Experiments. 

Experiment 133. Tests for Mercury. 

(a) What is a simple test for free mercury? 

(b) Recall or devise a test for combined mercury. Verify it. (Sug- 
gestion. See Exp. 115.) 

Experiment 134. Properties of Mercurous and Mercuric 
Compounds. Materials: Solutions of mercurous and mercuric 
nitrate ; hydrochloric acid, ammonium hydroxide. 

(a) Mercurous. Add a few drops of hydrochloric acid to a little 
mercurous nitrate solution. The white precipitate is mercurous chlo- 
ride. Note its insolubility in water and in dilute hydrochloric acid. 
Add a few drops of ammonium hydroxide. The black precipitate is 
mercurous ammonium chloride. Its formation is a delicate test for 
mercury in mercurous compounds. 

(b} Mercuric. Add a few drops of hydrochloric acid to a little 
mercuric nitrate solution. Compare the result with that in (a). Add 
a few drops of ammonium hydroxide, or enough to produce a decided 
change. Compare with (a). The precipitate is mercuric ammonium 

Exercises for Review. 

1. Describe the effect of heat on red oxide of mercury. What his- 
torical interest has this experiment ? 

2. What practical use has mercury ? 

3. What are amalgams ? 

4. What action has mercury upon gold ? 


Experiment 135. Properties of Aluminium. 

(a) Examine a piece of aluminium, and observe its properties. Has 
it any "spring" like brass ? Is it ductile, malleable, soft, hard, tough, 
brittle ? Will it melt in the Bunsen flame ? Try it. 

(b} Compare roughly the weight of a piece of sheet aluminium with 
a piece of pasteboard or glass having approximately the same volume. 
Estimate the specific gravity. Verify it by consulting a book. 

Experiment 136. Action of Aluminium with Acids and Alka- 
lies. Materials: Aluminium, sulphuric acid, hydrochloric acid, sodium 
hydroxide solution. 



() Add a small piece of aluminium to separate test tubes containing 
dilute sulphuric acid and concentrated hydrochloric acid. Warm, if 
necessary. Describe the action. Test the gas evolved. What com- 
pound is formed in each case ? 

(b} Add a small piece of aluminium to a test tube half full of dilute 
sodium hydroxide solution, and boil. Test any gas evolved. If only 
a little gas is liberated, attach a simple delivery tube and collect the 
gas over water. 

Other acids and alkalies act similarly. Draw a general conclusion 
from this experiment. 

Experiment 137. Preparation and Properties of Aluminium 
Hydroxide. Materials: Solutions of alum, sodium hydroxide, am- 
monium sulphide, and cochineal ; hydrochloric acid and ammonium 

(a) Add slowly a little sodium hydroxide solution to a test tube 
half full of alum solution. The gelatinous precipitate is aluminium 
hydroxide. Now add an excess of the alkali to one half, and dilute 
hydrochloric acid to the other. Describe the results. 

(b} Add a little solution of ammonium sulphide to a solution of alum. 
Describe the result. The precipitate is not a sulphide, but aluminium 
hydroxide, because aluminium forms no sulphide in the wet way. 

(c} Add a little alum solution to a dilute solution of cochineal, then 
add ammonium hydroxide. The colored product is called carmine lake. 
It belongs to a class of dyes formed by the combination of a vegetable 
dye and a metallic hydroxide, often aluminium hydroxide. 

Experiment 138. Tests for Aluminium. Materials for (c) : 
Aluminium sulphate ; cobaltous nitrate solution, blowpipe, charcoal. 

(a} What is a simple test for metallic aluminium ? 

(&) Recall or devise a test for combined aluminium. Verify it. 
How can aluminium compounds be distinguished from those of zinc? 

(c) Heat a little aluminium sulphate on charcoal in the blowpipe 
flame. Cool, and moisten with a drop of cobaltous nitrate solution. 
Heat again, and if the operation has been conducted properly, a blue 
residue will coat the charcoal. This color is characteristic of aluminium 
compounds. Compare this result with the action of other metallic 
compounds under similar circumstances. 

534 Experiments. 

Experiment 139. Preparation and Properties of Common 
Alum. Materials for (a) : Aluminium sulphate, potassium sulphate, 
evaporating dish. 

(a) Dissolve about 10 grams of aluminium sulphate in the least pos- 
sible amount of hot water. Dissolve 3 grams of potassium sulphate in 
the same way. Mix the clear, hot, saturated solutions in an evaporat- 
ing dish, and allow the solution to cool undisturbed. Crystals of 
potassium alum will be deposited. Remove the best ones; dry and 
examine. Describe them, giving color, luster, size, and crystal form. 

() Prove by actual tests that (i) they are a sulphate, and (2) they 
contain aluminium and water of crystallization. 


Experiment 140. Properties of Tin. Examine a piece of tin, and 
state its most obvious properties. Is it malleable, soft, hard, tough, 
brittle? Will it melt in the Bunsen flame? Try it. 

Experiment 141. Action of Tin with Acids. Materials: Tin, 
concentrated nitric and hydrochloric acids. 

(a) Put a small piece of tin in a test tube, cover with concentrated 
hydrochloric acid, add a little water, and heat in the hood. Heat 
gently at first, and when action begins regulate the heat accordingly. 
Most of the tin disappears, soluble stannous chloride being formed. 
Save this solution for Experiment 142. 

(<) Treat a small piece of tin with concentrated nitric acid in the 
hood. It is advisable to stand the test tube in the rack or in a bottle as 
soon as the action begins. The white, amorphous product is metastan- 
nic acid. How does the action of nitric acid on tin differ from and 
resemble its action on other metals, zinc, for example? 

Experiment 142. Tests for Tin. Materials for (c) : Solutions of 
mercuric chloride and stannous chloride. 

(a) What is a simple test for metallic tin? 

() Recall or repeat the action of tin when heated in a blowpipe 

(c) Add a few drops of mercuric chloride solution (poison) to a little 
of the stannous chloride solution prepared in Experiment 141 . The white 
precipitate is mercurous chloride. Add a little more stannous chloride 

Lead. 535 

solution and heat gently. The mercurous chloride is reduced finally to 
mercury, which appears as a grayish powder. 

Experiment 143. Deposition of Tin. 

Put a strip of zinc in a slightly acid solution of stannous chloride. 
Examine after a short time. The tin will be found adhering to the 
zinc as a grayish black deposit; sometimes bright scales are also seen. 
What has become of the zinc? 


Experiment 144. Properties of Lead. 

(a} Examine a piece of freshly cut lead and state its most obvious 
physical properties. 

() Estimate its specific gravity. Verify your estimate by consulting 
a book. 

(Y) Draw a piece of lead across a sheet of white paper, and describe 
the result. 

(d} Is lead easily melted? Try it. 


(1) What happens to lead when exposed to the air? 

(2) What properties adapt lead for its extensive use? 

(3) What is "black lead"? 

(4) Is there lead in a lead pencil? 

Experiment 145. Tests for Lead. Materials for (Y), (*/), (*) : 
Lead nitrate and potassium dichromate solutions, sulphuric acid, hydro- 
chloric acid. 

(a) Recall or repeat the reduction of lead oxide in the blowpipe 
flame. (See Exp. 79.) 

(d) Recall or repeat the action of hydrogen sulphide with the solu- 
tion of a lead compound. (See Exp. 95 ()) 

() Add dilute hydrochloric acid to a little lead nitrate solution until 
precipitation ceases. Note the insolubility of the lead chloride which is 
formed. Warm gently as long as any decided change occurs. Describe 
the action. This is characteristic of lead chloride and permits its 
separation from the chlorides of silver and of mercury (in the -ous 

53 6 Experiments. 

(d) Add dilute sulphuric acid to a little lead nitrate solution until 
precipitation ceases. The precipitate is lead sulphate. Observe its 
properties. Is it soluble in hot water ? Try it. 

(e) Repeat (</), using potassium dichromate solution instead of sul- 
phuric acid. The precipitate is lead chromate. Describe it, especially 
the color. 

Experiment 146. Deposition of Lead. 

Repeat Experiment 130 (b). 

Experiment 147. Properties of Lead Oxides. Materials: Lead 
monoxide, dioxide, and tetroxide, nitric acid. 

(a) Examine the three oxides and tabulate their most obvious physi- 
cal properties, stating the exact chemical name and formula and the 
popular name of each oxide. 

(b} Recall the experiment in which lead was heated in the oxidizing 
flame, especially the color of the coating. What oxide of lead is 
thereby formed ? 

(c) Warm a little lead tetroxide with dilute nitric acid. The solid 
product is lead dioxide. Describe it. 


(1) How might lead tetroxide be prepared ? 

(2) If lead tetroxide is heated strongly, lead monoxide is formed. 
What does this fact reveal about the stability of lead tetroxide ? 

(3) When lead dioxide and concentrated hydrochloric acid are mixed 
and heated, chlorine is evolved. Complete the equation 

PbO 2 + 4 HC1 = PbCl 2 + 2 H 2 O + 

How does this interaction resemble that of manganese dioxide and 
hydrochloric acid ? 

Experiment 148. Properties of Certain Lead Compounds. 

Materials: Lead nitrate, lead carbonate, galena, taper, mortar and 

(#) Put a crystal of lead nitrate in a test tube provided with a 
holder, hold in a horizontal position, and heat strongly in the upper 
part of the Bunsen flame. Describe the result. What is the most 
obvious product ? After most of this product has passed out of the 
test tube, thrust well into the test tube a taper with a spark on the 

Chromium. 537 

end. Describe the result. What other gas is present ? The solid 
product is lead monoxide. Summarize the behavior of lead nitrate 
when heated. 

(b) Examine lead carbonate, and state its most obvious properties. 
What is its common name ? Prove that it is a carbonate and con- 
tains lead. (Suggestion. Treat with hydrochloric acid.) 

(c) Examine a lump of galena, and state its most obvious proper- 
ties. Pulverize it in a mortar. What additional property is revealed ? 
Prove that it is a sulphide and contains lead. (Suggestion. See Exps. 
93 and 145.) 


Experiment 149. Tests for Chromium. Materials: Borax, 
chrome alum, potassium carbonate, potassium nitrate, acetic acid, nitric 
acid, sodium hydroxide solution, lead nitrate solution, potassium dichro- 
mate solution, platinum test wire, piece of porcelain, forceps. 

(a} Prepare a borax bead (see Exp. 102), touch it with a minute 
quantity of chrome alum, and heat in both the oxidizing and reducing 
flame. Describe the result. 

(#) Mix equal small quantities of potassium carbonate, potassium 
nitrate, and powdered chrome alum, place the mixture on a piece of 
porcelain, and hold it with the forceps in the upper Bunsen flame so 
that the mixture will fuse. A yellow mass, due to the presence 
of potassium chromate, results. If the color is not decided, dissolve 
the mass in water, add acetic acid, slowly at first, and boil to expel the 
carbon dioxide. Add a few drops of lead nitrate solution to a portion, 
and yellow lead chromate is precipitated. If the precipitate is white, 
it is lead carbonate, and shows that not all the potassium carbonate 
was decomposed, as intended. 

(c) Add lead nitrate solution to potassium dichromate solution. 
Name and describe the precipitate. Try the solubility of the precipitate 
in acetic acid, dilute nitric acid, and sodium hydroxide. Describe the 

Experiment 150. Properties of Chromates. Materials : Potas- 
sium chromate and dichromate, concentrated hydrochloric acid, 
potassium hydroxide solution. 

(#) Examine crystals of potassium chromate and dichromate, and 
state their characteristic properties. Make a dilute solution of each, 
and compare the colors. Save for (c) and (*/). 

53 8 Experiments. 

(b) State the properties of lead chromate. 

(V) Add a few drops of concentrated hydrochloric acid to the solution 
of potassium chromate prepared in (a), and observe the change in 
color. Describe it. Compare with the color of the potassium dichro- 
mate solution prepared in (a). Draw a conclusion. 

(d) Add potassium hydroxide solution to the solution of potassium 
dichromate prepared in (a) until a change of color is produced. De- 
scribe the color. Compare with the potassium chromate solution. 
Draw a conclusion. 

(e) The chromates are oxidizing agents. Add a few drops of con- 
centrated hydrochloric acid to powdered potassium chromate and dichro- 
mate in separate test tubes. Chlorine is evolved. Where did it come 
from ? By what general chemical change ? 

Experiment 151. Reduction of Chromates to Chromic Com- 
pounds. Materials: Potassium dichromate solution, concentrated 
hydrochloric acid, alcohol. 

Add to a few cubic centimeters of potassium dichromate solution a 
little concentrated hydrochloric acid and a few drops of alcohol. Warm 
gently. Two important changes occur. The chromate is reduced to 
chromic chloride which colors the solution green ; the alcohol is oxi- 
dized to aldehyde, which is detected by its peculiar odor. 

Experiment 152. Preparation and Properties of Chromic 
Hydroxide. Materials: Ammonium sulphide, solutions of sodium 
hydroxide and chrome alum. 

(a) Add a little sodium hydroxide solution to a solution of chrome 
alum. The precipitate is chromic hydroxide. Describe it. Add an 
excess of sodium hydroxide solution, and shake. Describe the result. 
Boil, and state the result. 

(d) Add a little, and then an excess, of ammonium sulphide to a 
solution of chrome alum. Compare the result with that in (a) . Does 
chromium form a sulphide ? 


How can chromic hydroxide be distinguished from aluminium 
hydroxide ? 

Manganese. 539 

Experiment 153. Properties of Chrome Alum. 

() Examine chrome alum and state its most obvious physical 

() Prove that chrome alum is a sulphate, and that it contains 
chromium and water of crystallization. 


Experiment 154. Tests for Manganese. Materials for (b) and 
(c) : Manganese dioxide, potassium carbonate, potassium nitrate, am- 
monium sulphide, manganese sulphate solution, hydrochloric acid, 
acetic acid, ammonium hydroxide. 

(a) Subject a minute quantity of manganese dioxide to the borax 
bead test, and note the color of the bead after heating in each flame. 
(See Exp. 102.) 

(b) Fuse, on a piece of porcelain, a little manganese dioxide 
mixed with potassium carbonate and potassium nitrate. The green 
mass is a test for manganese. It is due to the presence of potassium 

(c) Add ammonium sulphide to manganese sulphate solution. The 
flesh-colored precipitate is manganese sulphide. Compare with other 
sulphides as to color (see Exp. 95). Divide it into two parts. Add 
hydrochloric acid to one, and acetic acid to the other, then add an 
excess of ammonium hydroxide to each. Draw a conclusion regarding 
the solubility of manganese sulphide. 

Experiment 155. Oxidation with Potassium Permanganate. 

Materials: Potassium permanganate, sulphuric acid, ferrous sulphate, 
filter paper. 

(a) Add a few drops of sulphuric acid to a weak solution of fresh 
ferrous sulphate ; then add, drop by drop, a dilute solution of potas- 
sium permanganate. Its color is changed, owing to the loss of oxygen ; 
the latter converts the ferrous to ferric sulphate. The decomposi- 
tion of the permanganate also causes the formation of potassium and 
manganese sulphates. 

(b) Pour a solution of potassium permanganate upon a piece of filter 
paper. Describe and explain the result. 

540 Experiments. 

Exercises for Review. 

1. Describe manganese dioxide. Name five elements in whose prepa- 
ration manganese dioxide is used. Is manganese dioxide an oxidizing 
agent ? 

2. Describe potassium permanganate. What can be said of its solu- 
bility in water ? In what previous experiment has it been used ? 

3. What is the formula of potassium permanganate ? Does the 
formula suggest its oxidizing power ? 


Experiment 156. Properties of Iron. Materials: Cast and 
wrought iron, steel, magnet, iron wire, iron powder. 

(a) Examine cast iron, wrought iron, and steel, and state their most 
obvious physical properties 

() Try the action of a magnet on each. Describe the result. 

(c) Drop a pinch of iron powder into the Bunsen flame. Hold a 
piece of fine iron wire in the flame. Describe the results, and draw a 

Experiment 157. Properties of Ferrous Compounds. Mate- 
rials: Iron powder (or filings), hydrochloric acid, solutions of sodium 
hydroxide, potassium ferricyanide, potassium thiocyanate, potassium 

Put a few grams of iron powder in a test tube, add about 10 cubic 
centimeters of dilute hydrochloric acid, and warm gently ; ferrous chlo- 
ride is formed (in solution). Proceed as follows : (i) Pour a little into a 
test tube one third full of sodium hydroxide solution. The precipitate is 
ferrous hydroxide. Watch the changes in color. To what are the 
changes due ? (2) Add a second portion to potassium ferricyanide 
solution. The precipitate is ferrous ferricyanide. Describe it. (3) Add 
a third portion to potassium thiocyanate solution. If ferric salts are 
absent, no change results. (4) Add a fourth portion to potassium 
ferrocyanide solution. The precipitate is ferrous ferrocyanide. Describe 
it. Tabulate the results as described in the next experiment. 

Experiment 158. Properties of Ferric Compounds. Materials : 
Ferric chloride solution and the solutions used in Experiment 157. 

Iron. 541 

To a little ferric chloride solution add (i) sodium hydroxide solu- 
tion. The precipitate is ferric hydroxide. Describe it. Add to ferric 
chloride solution {2) a little solution of potassium ferricyanide. Com- 
pare the negative result with (2) in Experiment 157. Add as above (3) 
a little solution of potassium thiocyanate. The rich wine-red coloration 
is caused by the soluble ferric thiocyanate. This test distinguishes ferric 
from ferrous compounds. Add as above (4) a little solution of potas- 
sium ferrocyanide. The precipitate is ferric ferrocyanide. Describe it. 
Tabulate the results of these two experiments, showing the behavior 
of ferrous and ferric compounds under the same conditions. 

Experiment 159. Reduction of Ferric Compounds. Mate- 
rials: Ferric chloride solution, zinc, hydrochloric acid. 

Put a piece of zinc in ferric chloride solution made slightly acid by 
hydrochloric acid. The nascent hydrogen reduces the ferric to ferrous 
chloride. After the operation has proceeded for about fifteen minutes, 
test a portion of the liquid for a ferrous and a ferric compound by Ex- 
periments 157 (2) and 158 (3). If the tests are not conclusive, continue 
the reduction and test again. Describe the result. 

Experiment 160. Oxidation of Ferrous Compounds. Mate- 
rials: Ferrous sulphate, hydrochloric acid, potassium chlorate, nitric 

(a) To a solution of fresh or freshly washed ferrous sulphate add a 
little hydrochloric acid, warm gently, and then add a few crystals of 
potassium chlorate. After heating a short time, test portions of the 
liquid for a ferric and a ferrous compound. 

(b) Add 10 cubic centimeters of concentrated nitric acid, drop by 
drop, to a hot solution of ferrous sulphate to which a little sulphuric 
acid has been added, and boil. Test portions of the liquid for a ferric 
and a ferrous compound as in Experiment 159. 

(c} Recall a third illustration of the oxidation of a ferrous to a ferric 
compound. Describe it briefly. 

Experiment 161. Properties of Certain Iron Compounds. Ma- 
terials : Ferrous sulphate, hematite, li'monite, magnetite, pyrite, siderite. 

(a} Examine a crystal of ferrous sulphate, and state its most obvious 
properties. Heat it gently in a test tube inclined mouth downward. 
Describe the result. Test this crystal for ferrous and ferric compounds 


as in Experiment 159. State and explain the result. What is the 
common name of ferrous sulphate? 

() Examine hematite, limonite, and magnetite, and state their prop- 
erties. Draw the first two across a rough sheet of paper or a piece of 
ground glass, and describe the "streak " made by each. What is the 
formula of each compound? Prepare a hydrochloric acid solution of 
each and test for iron. State the result. 

(V). Examine pyrite, and state its properties. It is iron disulphide. 
What is its formula? For what is it used? For what is it sometimes 
mistaken ? 

(d) Examine siderite, and state its properties. It is ferrous carbon- 
ate. What is its formula? Test a powdered specimen for a carbonate 
and for iron. State the result. 

Exercises for Review. 

1. What happens when iron is (a) treated with acids, and (b) heated 
with sulphur? 

2. Describe ferrous sulphide. What are formed by its interaction 
with warm hydrochloric acid? 

3. What happens to iron when it is placed in copper sulphate 
solution ? 


Experiment 162. Test for Nickel. 

To a solution of nickel chloride add sodium hydroxide to alkaline 
reaction. The precipitate is nickelous hydroxide. Describe it. 

Experiment 163. Test for Cobalt. Repeat Experiment 102 (a). 


Experiment 164. Composition of Organic Compounds, Ma- 
terials: Turpentine, alcohol, camphor, kerosene, sugar, starch, flour, 
wood, paper, hair, candle, taper, gelatine, mustard, silver coin, red 
litmus paper, soda lime, porcelain dish, kerosene lamp, two bottles. 

(a) Carbon, (i) Recall or repeat the experiments which showed 
that carbon is a constituent of wood, cotton, starch, sugar, illuminating 
gas and candle wax. (2) Heat a few drops of turpentine in a porce- 
lain dish, and then set fire to it. Does it contain carbon? Hold a 
bottle over the flame long enough to collect any product, and then test 

Organic Compounds. 543 

the contents for carbon dioxide with limewater ; does the observation 
verify the previous conclusion? (3) Repeat (2) with alcohol. Does 
alcohol contain carbon? (4) Burn a small lump of camphor in a porce- 
lain dish or on a block of wood. Does it contain carbon? (5) Hold 
a bottle over a burning kerosene lamp long enough to collect any prod- 
uct, and test as in (2). Does kerosene contain carbon? 

(b) Hydrogen, (i) Set fire to a few drops of the following liquids 
in a porcelain dish, and hold over each flame a cold dry bottle long 
enough to allow the condensation of the water vapor, which is always 
one product of the combustion of organic compounds which contain 
hydrogen : alcohol, turpentine, kerosene. (2) Heat in separate test 
tubes the following dry solids, and if they contain hydrogen, a little 
water vapor will condense on the upper part of the test tube : sugar, 
starch, flour, wood, paper, hair. (3) Hold a cold, dry bottle for a few 
seconds over a burning kerosene lam